403

J. Electroanal. Chem., 252 (1988) 403-424 Elsevier Sequoia S.A., Lausanne - Printed in The Netherlands

The adsorption of stereoisomers at a Hg electrode

The effect of concentration and temperature

Fernando Silva and Morn6 Rodrigues

Departamento de Quimica, Faculdade de Ci&cias, Unroersrdade do Porte, 4000-Port0 (Portugal)

(Received 18 April 1988; in revised form 13 May 1988)

ABSTRACT

The adsorption of mannitol and sorbitol was studied as a function of the bulk concentration. The degree of adsorption follows the order of the solubilities of the compounds. This result is explained by overriding solute-solvent interactions since the vertical interactions of both polyalcohols with the electrode surface are similar. The effect of temperature on the differential capacity and the derived entropy of formation of the interface corroborates the idea that a water structure different from that in the bulk is responsible for the discrimination in adsorption of the small structural differences between the two solutes. Some results obtained for another isomer of sorbitol, dulcitol, are in qualitative agreement with the model proposed.

INTRODUCTION

Specific interactions between polyols and water have been inferred from dielec- tric and NMR relaxation [l], partial molar volumes and compressibilities [2,3], apparent molal heat capacities [4] and aqueous heat capacities [5].

Mannitol and dulcitol are biologically important polyalcohols, whose structures differ from that of sorbitol only by the position of one hydroxyl group (Fig. 1).

Preliminary studies on the adsorption of mannitol and sorbitol on Hg made by Peat and Shannon [6] indicated substantial differences in their adsorptive behaviour despite the similarity of their structures. However, it was not possible to establish whether the difference in adsorption was due to specific metal-adsorbate interac- tions of solvation.

In an attempt to understand the adsorption of these stereoisomers, a systematic study of the adsorption of mannitol and sorbitol on Hg from aqueous solutions was undertaken. Some results obtained for dulcitol, which is another isomer of sorbitol, are also presented.

0022-0728/88/$03.50 0 1988 Elsevier Sequoia S.A.

404

i

- H20H

i

-HZOH C- H20H

I HO- C-H H-C-OH H -C-OH

I I I HO-C-H HO-C-H HO-C-H

I I I H-C-OH H -C-OH HO-C-H

I I I H-C-OH H-C-OH H- C-OH

I

I I C - H20H C- H2OH C-H2OH

D-mannitol sorbitol

Fig. 1. Structures of manmtol, sorbitol and dulcitol.

dulcltol

Account is given here of the characteristics of the adsorption of these stereoiso- mers in aqueous NaCl solutions, in a range of concentrations. Despite the fact that chloride ion might be adsorbed, the use of NaCl as supporting electrolyte is justified because there are thermodynamic data for such solutions [7] and because it is a component of biological fluids. Furthermore, it has been proposed that NaCl is effective in differentiating the structural effects of stereoisomeric molecules [4].

EXPERIMENTAL

NaCl, mannitol and dulcitol were recrystallized twice from ultrapure water. Sorbitol was very difficult to recrystallize from water and purification was achieved by recrystallization from 95% ethanol. All solutions were prepared with water from a Millipore super Q system. For mannitol + NaCl + water solutions, the activity of NaCl was kept constant by adjusting the concentration of NaCl according to the thermodynamic data available [7].

The cell contained a compartment for the reference electrode (calomel with 1 mol dm-3 NaCl solution to avoid liquid junctions) and was water-jacketed. The solu- tions and electrodes were thermostated using the circulation from a thermostatic bath (+0.2O C). Prior to any measurements, oxygen was removed by bubbling purified nitrogen through the solutions.

Measurements of the differential capacity were made using a lock in amplifier technique together with an automatic data acquisition system.

The total capacity of a mercyry drop can be obtained from the in-phase, V, and out-of-phase, V4, components of the ac potential from

(1)

405

drop tbme

Fig. 2. Potential-time programme used in the measurements of the differential capacity.

where K is a constant which can be obtained with an appropriate dummy cell. However, there are difficulties in applying eqn. (1) to a Hg drop due to the fact that the area changes with time. To overcome such difficulties and therefore to increase the precision, the differential capacity was obtained from a computer fit of the measurements of C, during the drop life. The differential capacity C is related to the total capacity C, through

C, = C(6m7r,‘p)2’3)t2’3 - CA, (2)

where p is the density of mercury, t is the time after drop birth, m is the mercury mass flow rate and A, is the shielded area of the capillary.

From a previous study [8], it was found that the best linearity in the whole range of potentials employed was obtained by taking 75 measurements at intervals of 50 ms, starting 3 s after drop birth. Correlation coefficients of the straight lines adjusted to the (C,, t) data were always above 0.999 and the values of A, obtained from the intercepts, although varying with potential, were in the range 2.8-3.7 x 10m4 cm*, which is close to the value estimated from the diameter of the capillary bore (3.2 x lop4 cm*). For a fully automatic procedure a staircase wave of potential synchronized with the drop birth was employed. Figure 2 illustrates the time programme used.

Drop time electrocapillary curves were obtained with a home-made timer, the details of which are given elsewhere [9]. Calculation of the values of surface tension at the electrocapillary maximum was made after calibration with a 1 mol dmP3 NaCl solution and using the method described in ref. 10.

RESULTS

Effect of concentration

Figures 3 and 4 illustrate the dependence of the differential capacity on the applied potential for different concentrations of each solute.

406

40

30

298 K

N&l 1 mol dm-3

Mannltol 0 05 0 10

-E/ “(WE)

I J

Fig. 3. Differential capacity-potential curves of Hg/l mol drnm3 NaCl+mannitol as a function of the potential. The concentration of mannitol is indicated next to each curve.

The curves do not show the typical organic adsorption-desorption peaks; only sorbitol displays small maxima at high negative potentials. The absence of well- defined adsorption/desorption peaks is characteristic of a weak interaction between molecules in the adsorbed layer [ll]. The results obtained for manmtol and sorbitol are similar to those reported by Parsons and Peat [12] and Peat and Shannon [6] in fluoride ion solutions. This seems to indicate that the adsorption behaviour of the two polyhydroxy compounds is similar in the presence of either chloride or fluoride anions. Figure 5 shows the dependence of the differential capacity of dulcitol, another isomer of sorbitol, on the concentration and potential. A comparison of the C(E) curves for the three compounds made in Fig. 6 indicates that the capacity in the presence of sorbitol is lower than in the presence of either of its stereoisomers, mannitol and dulcitol. This may be taken as indicative that sorbitol adsorbs to a greater extent than either of the other two compounds. At identical concentrations, the C(E) curves of dulcitol on Hg are very similar to those of manmtol, although their solubilities in water are very different; dulcitol is much less soluble than mannitol.

407

40

30

298 K

N&i 1 moi dm-

Fig. 4. Differential capacity-potential curves of Hg/l mot dme3 NaCl+sorbitol as a function of the potential. The concentration of sorbitol is indicated next to each curve.

Integration of the C(E) curves was performed numerically [9] using as integra- tion constants the values of the pzc obtained from differentiation of drop-time electrocapillary curves. Integration of the C(E) curves yielded charge-potential curves with a common point, the coordinates of which differ for each compound and are given in Table 1.

The charge at the common point, u,_, corresponds to m~mum adsorption of each solute and the values become more positive going from sorbitol to mannitol and to dulcitol. The values of amax are slightly more positive than those obtained by Peat and Shannon [6] which may be due to the presence of chloride ion at the interface.

Figure 7 shows the dependence of the pzc on the concentration for mannitol (a) and sorbitol (b). The shift of the pzc towards more negative values with an increase in the concentration of each solute is similar to the results observed by Peat and Shannon and points to an increase in adsorption of each polyol with increasing concentration. The similarity of the curves may be interpreted as suggesting that the dipole moments of the adsorbed molecules are not very different and are consistent

408

40

298 K

- E/ “(SCE)

15 10 05 -I

Fig. 5. Differential capacity-potential curves of Hg/l mol dme3 NaCl +dulcitol as a function of the potential. The concentration of dulcitol is indicated next to each curve.

with the absence of any noticeable differences between the C(E) curves in the presence of chloride or fluoride anions.

Further integration of the a-E curves was done using as integration constant the value of y,,, obtained from the drop time measurements. The values of the surface pressure at constant charge density were used to test for congruency of the isotherm. The surface pressure-log a,,,,,,,, and log c,,,~~~~, curves were superimposable on top of the curve corresponding to the charge of maximum adsorption by translation along the abscissae, in the range of surface charge densities - 10 to + 6 PC cme2. The composite curves are shown in Fig. 8; the experimental scatter is of the order of +0.5 mN m-‘. This indicates congruency of the isotherm and made it easier to fit an isotherm to the data. The data could not be described by one-parameter isotherms and best fits were obtained with a Frumkim isotherm using the parame- ters given in Table 2. Thermodynamic data are not available for the ternary system sorbitol + NaCl + water but a comparison with the mannitol system suggests that the effect of the medium on the activity of the supporting electrolyte is negligible in the concentration range studied. Although there are differences in the properties of

409

40 Mannltol 0 1 mol dm3

Dulcltol 0 1

Sorbliol 0 1

-E/V vsSCE

15 10 05

Fig. 6. Comparison of the C(E) curves for mannitol, sorbitol and dulcitol.

sorbitol and mannitol solutions, the assumption was made that the activity of sorbitol could be approximated by the concentration. To evaluate this procedure, a test was made with the data for mannitol, carrying out the analysis using concentra-

tions and activities; the difference in the results obtained was found to be very

TABLE 1

Summary of the adsorption parameters

Compound %ax / ElII,, / pC cm-2 V (SCE)

a 10” r,/ AG”/ mol cm-’ kJ mol-’

Sorbitol +1.0 - 0.490 - 0.3 32 - 12.4 Dulcitol + 2.0 - 0.473 Mannitol + 3.0 - 0.460 + 0.3 25 - 15.4 Xylose a + 2.6 0.382 b - - 0.25 33 - 14.5 Sucrose ’ 0 + 2.0 24 - 22.1

a Data taken from ref. 12. b Value referred to a 0.795 M reference electrode.

3 c imol.dri3

0.1 0.2 0.3 0.4 0.5 0.6 1 I I I

Fig. 7. Dependence of pzc on the solute concentration. (a) Mannitol; (b) sorbitol.

small, both in the fitting and in the differentiation processes. This may be taken as indicating that the errors introduced by the approximation are likely to be of the same order of magnitude as those of the data or introduced by the numerical treatment.

The values of the saturation coverage, I’,, are not very different for mannitol and sorbitol, indicating that the area occupied by each molecule at the surface is similar for each compound. However, there are noticeable differences in the values of the standard Gibbs energy of adsorption, AG”, and particularly in the value of the interaction parameter a which, although small, has opposite signs for the two polyols.

The dependence of the relative surface excesses on u, obtained from numerical differentiation of the surface pressure (p = tb - 5, vs. log c, where 5 = y + aE, is shown in Figs. 9 and 10. The striking feature is the non-symmetrical curves obtained for mannitol. The deviations of positive charge densities may be the result of coadsorption of chloride ion. In accordance with the C(E) curves, the data indicate that sorbitol adsorbs more strongly than mannitol, which is unexpectedly the same order as that of their solubilities.

411

t

15

lo3 4 IN m-’

0 -3 -2 -1 0 L 1 8

Fig. 8. Composite surface pressure for mannitol (a) and sorbitol (b) adsorbed on mercury in contact with 1 mol dmV3 NaCl at 25 o C. The tines were calculated from the Frumkin isotherm using the parameters

given m Table 1.

The standard energy of adsorption was calculated from the adsorption coeffi- cient, log /?:

log j3 = log(l/c,.b) - bGO/RT (3)

where c,,~ is the bulk solvent concentration. The magnitude of the AGo values is similar to that obtained for the adsorption

of xylose [13] and points to physical adsorption of mannitol and sorbitol without the existence of specific interactions of the molecules with mercury. The variation of the standard Gibbs energy of adsorption with charge density was obtained from the horizontal displacements [14] to obtain the composite surface pressure-log asolUte curves. In both cases, a nearly quadratic dependence of AG” on u characteristic of

TABLE 2

Dependence of the temperature invariant point in the differential capacity on the concentrations of mannitol, sorbitol and dulcitol, in aqueous 1 mol dmS3 NaCl

Concentration /mol drne3

0.025

0.075 0.1 0.3

0.5

udc/dWJ/Pc cm-2

Mannitol

- 4.0 + 3.1

+0.5

Sorbitol

+ 5.0 + 4.0

+2,1

Dulcitol

+6.8

+ 6.0 + 5.4

412

Id' r Mannltol imoi cme2

/ a/pC cm-*

0 10 20

Fig. 9. Relative surface excesses of mannitol as a function of the surface charge density.

organic adsorption was found although small deviations were observed at positive charge densities, particularly for mannitol.

The values of the interaction parameters are small, suggesting weak interactions at the surface which is consistent with the absence of adsorption/desorption peaks on the C(E) curves. The sign of the values of the interaction parameter has been taken as evidence of the type of interaction predominant at the surface [15,16]. For mannitol, the positive value of a indicates that the major interactions are repulsive, while the negative value obtained for sorbitol points to net attractive interactions in the interfacial solution. The existence of such opposite net effects may be interpre- ted by suggesting that the solvent-solute interactions in the interphase are different for each compound. This presumably must arise to some extent because of the difference in the structure of water at the interphase and in the bulk. The difference in behaviour between xylose and the more soluble sucrose was also interpreted in terms of a specific water structure at the surface [13]. In order to explain the difference in adsorption at the air/water and Hg/water interfaces of mannitol, sucrose and a few other compounds, the existence of different structures of water at those interfaces has also been proposed [ll]. Furthermore, results obtained for the adsorption of D-ribose and 2-deoxy-ribose [17] illustrated that small changes in the structure of the molecules can cause marked differences in the adsorption be-

413

050 pj

0 30

020

010

005

Fig. 10. Relative surface excesses of sorbitol as a function of the surface charge density.

haviour. The structural differences between sorbitol and its isomers mannitol and dulcitol are also very small and it seems that the structure of water at the surface is able to discriminate between some of those differences.

From studies in solution the results support the view that the differences in behaviour between mannitol and sorbitol are not due to long-range solute-solute interactions but to short-range solute-solvent effects [4,5]. It has been suggested [18] that sorbitol disrupts the solvent structure to a greater extent than mannitol because it is unable to assume a conformation compatible with the solvent H-bond network. The positions of the hydroxyl groups in the mannitol molecule are such that marmitol is allowed to enter the hydrogen network of bulk water with minimum disturbance. The surface disturbs the hydrogen network of water molecules [19] and, consequently, the molecules of mannitol will find an environment not favourable to establish hydrogen bonds, with the solvent molecules giving rise to a net repulsive interaction which diminishes the amount adsorbed. Similarly, the unadaptability of the dulcitol molecule to the surface water structure may explain its behaviour comparatively to that of mannitol or sorbitol, as conveyed by the differential

capacity curves. In contrast, the molecules of sorbitol which do not adapt to the solution H-bond network may find the surface water less unfavourable and, therefore, accumulate at the interphase more easily than mannitol. In conclusion, it may be suggested that the nature of the solvent-solute interactions near the surface,

414

C / pF cm-2

Fig. 11. Differential capacity of Hg/l mol dm-3 NaCl+O.3 mol dme3 mannitol as a function of the charge and temperature.

mediated by a different structure of the solvent, is responsible for the difference in adsorption between mannitol and sorbitol.

Effect of temperature

Measurements of the differential capacity of the solutions of mannitol, sorbitol and dulcitol were made at different temperatures ranging from 5 to 45OC. The results obtained were integrated using the pzc obtained at each temperature and typical results are shown in Figs. 11, 12 and 13. The effect of temperature on the differential capacity is qualitatively identical to what has been observed for other compounds [20,21], namely that there are two domains of polarization where the temperature coefficient of the differential capacity has opposite sign. The isosbestic point that separates these two domains has coordinates which depend on both the nature and the concentration of the particular solute.

The isosbestic points occur at positive charge densities and the values are given in Table 2 for each solution.

415

-2 a/NC cm -10 0 10

Fig. 12. Differential capacity of Hg/l mol drnm3 NaCl +O.l mol dm-3 sorbitol as a function of the charge and temperature.

The dependence of the values of the temperature coefficient of the inverse of the differential capacity on the charge density is given in Figs. 14 and 15. The data show marked differences between the three compounds, both in the absolute values of the temperature coefficients and in the effect of concentration. The invariance of the charge at the maximum in d(l/C)/dT in the mannitol solutions is noticeable: for sorbitol the value of u for the maximum shifts towards more positive values with an increase in concentration while for mannitol the maximum in d(l/C)/dT is practically invariant with the concentration. The behaviour observed for dulcitol is intermediate between that of sorbitol and mannitol.

The calculation of the entropy of formation of the interphase was carried out according to established routes [20-221. The first integration yields the dependence of the generalized electrochemical potential l on the charge density and requires as integration constant the temperature coefficient of the inner potential drop at the uncharged Hg/solution interface (dA+(Hg-sol)/dT). This value may be obtained from the experimental temperature coefficient of the pzc and from the temperature coefficient of the reference electrode used in the isothermal cell according to

[dA+(Hg-sol)/dT] o__o = d(pzc)/dT + d E,,,/dT (4)

416

C /pF cm-2

O/UC cm-2

Fig. 13. Differential capacity of Hg/l mol dmm3 NaCl+0.075 mol dm-’ dulcitol as a function of the charge and temperature.

The temperature coefficient of the reference electrode, d E,,/dT, was taken from the literature [23]. Values of d(pzc)/dT calculated from the experimental data are given in Table 3; all the temperature coefficients are positive and no systematic variation with concentration was found. The second integration gives the entropy of formation, S *(a), according to

S*(a) - S*(q) = -LO /O(d(l/C)/dT),,O + (d(pzc)/dT+ dE,,/dT) da 9 1 (5)

TABLE 3

Temperature coefficients of the point of zero charge for mannitol, sorbitol and dulcitol solutions

Mannitol Sorbitol Dulcitol

d(pzc)/dT/mV K-’ + 0.43(8) + 0.41(2) + 0.47(5)

417

105x[d(l/C)/dT] / pF-‘K ‘cm2

Fig. 14. Dependence of d(l/C)/dT on the surface charge density of mannitol and dulcitol in 1 mol dmm3 NaCl solutions at 298 K.

O/PC cm 2

-10 0 10

The integration constant should be the value of S* at a charge of reference. However, knowledge of the variation of the absolute values of S * is not essential to this study; therefore the integration constant was arbitrarily, but conveniently, made zero for all solutions at (I = 0, i.e. S *(u = 0) = 0.

For solutions of mannitol, the C(a) curves apparently merge with that of the base electrolyte at charges more negative than -14 PC cmP2. This experimental result may be taken as indicative that the composition of the interface is no longer dependent on the concentration of mannitol in solution and therefore that S* curves should also merge at values of u more negative than - 14 PC cmP2. The forced fit of the S*(u) curves for mannitol is given in Fig. 16. It must be emphasized that such a fit affects only the vertical displacements of each curve. The important aspect of all the curves is that each shows a maximum, the position of which is systematically dependent on the composition of the solution and, therefore. of the interface. The effect of adsorption of mannitol is to decrease the entropy of formation of the interphase. The decrease of S* with lYmannlto, at constant solution

418

105xId(liC)ldT] ,/ gF ‘K-‘cm2

SorbMl

Fig. 15. Dependence of d(l/C)/dT on the surface change density of sorbitol in 1 mol dm-’ NaCl solutions at 298 K.

composition is evidence that there is a net negative entropy in the adsorption process. The lateral displacement of the maximum of S * may be explained by such a negative value of the entropy of adsorption and may also be suggested as evidence of the reorganization of the complex interface [21].

The dependence of the entropy of formation S*(a) on the concentration of sorbitol is given in Fig. 17. The curves are also of quasi-parabolic shape but no attempt was made to adjust them to a common region because the capacity curves are very different at large negative values of u. Even with the curves referred to the integration constant chosen it is noticeable that there is a shift of the maxima in the curves towards more negative values of u. However, this displacement is much smaller than that observed for mannitol, which suggests a lower value of the entropy of adsorption. The maximum of S * for the dulcitol interface occurs at charge densities more negative than is observed for sorbitol (Fig. 18).

A residual entropy &S * = S,* - S c has been defined [20] which allows the

419

Fig. 16. S* as a function of D and of the bulk concentration of mannitol. The curves were forced to fit that of the base electrolyte at negative charge densities.

determination of the effect of adsorbed molecules on the entropy of formation of the interphase and hence on the entropy of adsorption of the solutes.

An attempt was made to calculate AS* using the equation

AS*(u) = /( [ ~“@(W,)/d’%b.o do - j-,"(d(l/Co)/dT),o~. da] + K do]

(6)

where K is the difference between the temperature coefficients of the points of zero charge for the base electrolyte and for the organic solution measured against the same reference electrode.

The results obtained are given in Fig. 19 and the falling AS *(a) curves reflect the negative entropy of adsorption of each solute. Both curves have extrema at negative charges close to the values of the charge of maximum adsorption and closely resemble the form of the isotherms. This suggests that the decrease in AS* is predominantly due to a change in th@ relative surface excess of mannitol and sorbitol. It remains that the falling AS* curves reflect the change in the interfacial composition which is further reinforced by the dependence of AS* on the surface

420

Fig. 17. S* as a function of o for the concentrations of sorbitol indicated. The integration constants

were arbitrarily made zero at (I = 0.

excess of mannitol or sorbitol as shown in Fig. 20 and 21. According to the definition of AS* [21,22] and since the surface excesses are closely related to surface concentrations, it follows that

d(AS*)/dT,=(S;;-So)-n(S;zo-S,zo) (7)

Fig. 18. S* as a function of 0 for the concentrations of dulcitol indicated. The integration constants

were arbitrarily made zero at 0 = 0.

421

AS’

1

I 2.10-J K-1 cm-2

Fig. 19. AS* as a function of s for (a) 0.5 mol dmM3 mannitol and (b) 0.5 mol dmm3 sorbitol.

where AS, = (S; - So) is the variation of the partial molar entropy of the organic species with adsorption, ASHZO = (S&*, - SHzO ) is the corresponding value for the water molecules, and n is the number of water molecules displaced or replaced by adsorption or desorption of a molecule of solute.

The plots in Figs 20 and 21 are linear, suggesting that the partial molar entropy of adsorption of each solute is the predominant term in eqn. (7) which is negative and overrides the charge dependence of the other term, known to exist from previous studies [20]. The deviations of linearity observed for mannitol at positive charge densities are, as suggested earlier, connected with chloride ion adsorption. No special effect is observed which could be attributed to any difference in the partial molar entropy of chloride ion in each solution or to different contributions of the diffuse part of the double layer.

The lines for mannitol and sorbitol have similar slopes for all the concentrations, and the difference in the absolute values of AS *, even at constant r, results solely

422

Fig. 20. AS* as a function of r at two different bulk concentrations of mannitol in 1 mol drnm3 NaCl

at 298 K.

Fig. 21. AS* as a function of r at two different bulk concentrations of sorbitol in 1 mol drne3 NaCl at 298 K.

from the choice of the integration constant, which again was made arbitrarily zero for all concentrations at u = 0.

The values of a(AS */al?,) obtained from Figs. 20 and 21 are - 16 J K-r mol-’ for sorbitol and -45 J K-’ mol-’ for mannitol. As expected, the values are negative and contain an unknown contribution due to water displacement. Never- theless, the difference between mannitol and sorbitol may be interpreted as being the result of a more rigid distribution of mannitol molecules at the surface compared with that of sorbitol. The large value of the slope of the lines in Fig. 20 explains the shift towards more negative values of u of S* as observed in Fig. 16. Although S*(a) and AS * (I’) for mannitol are similar to those observed in the case of strongly adsorbing species such as thiourea [21] and iodide ion [24], where similar plots were also found to be linear, the effect of adsorption of mannitol on the differential capacity curves is very different. This contrasting behaviour may be the result of different overriding interactions: with the solvent in the case of mannitol and with the metal in the case of iodide or thiourea.

These results may be interpreted by suggesting that adsorption of mannitol yields a structured solute-solvent arrangement at the surface as the result of a compromise for its unadaptability to the solvent structure and its ability to form H-bonds. Such a situation would probably explain the net repulsive character of the interaction coefficient in the isotherm and, consequently, the amount of mannitol adsorbed. The adsorption of sorbitol cause a smaller variation of entropy because it disturbs the solvent arrangement at the surface less.

423

CONCLUSIONS

The effect of the concentration on the adsorption of mannitol and sorbitol has allowed us to obtain the isotherms. The parameters of the isotherms may interpret the differences in adsorption as the result of overriding solute-solvent effects. The difference in the interactions mannitol-interfacial water and sorbitol-interfacial water is the result of a specific structure of water at the surface which allows small structural changes to be enhanced in the adsorption behaviour. A comparison with dulcitol further reinforces that it is the solvent-solute interaction at the surface which allows the discrimination of the small structural changes to be observed with the adsorption process.

The effect of temperature, in particular the entropy of formation of the interface and its dependence on the composition, corroborates such a model. Large negative values for the variation of entropy with adsorption for mannitol may be the result of different interactions with the solvent at the surface.

ACKNOWLEDGEMENTS

F. Silva gratefully acknowledges the financial support of the Gulbenkian Foun- dation and INIC-UP Linha 3. The authors are indebted to Dr. D. Schiffrin for reading the manuscripts and for valuable discussions.

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424

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