Colloids and Surfaces, 51 (1990) 49-60 49 Elsevier Science Publishers B.V., Amsterdam

The adsorption of mannitol at the mercury/aqueous sodium fluoride interphase

Roger Parsons and Robert Peat 1

Department of Chemistry, University of Southampton, Southampton S09 5NH (United Kingdom)

(Received 5 April 1990; accepted 11 May 1990)

Abstract

The adsorption of mannitol was studied by measuring the interfacial tension and the differential capacitance of a mercury/aqueous electrolyte interphase. The results are compared with those obtained in the presence of chloride adsorption and with similar results for sucrose. Chloride adsorption has a minor but significant effect on the adsorption of these molecules and its nature suggests that classical models of non-electrolyte adsorption cannot be applied to these systems.

INTRODUCTION

The interest in the study of the adsorption of sugars, from both the funda- mental view point, as well as for its relation to biological systems has already been emphasized [ 1,2 ]. Detailed studies of sucrose [ 3,4 ] and of xylose [ 5 ] have been published. More recently Silva and Rodrigues [6] have followed up sug- gestions about the remarkable differences in the adsorption of stereoisomers [ 1,2,7 ] by studying the behaviour of mannitol, sorbitol and dulcitol. This work was done in an aqueous NaC1 base solution and there is a surprisingly large difference in the capacity curves obtained for mannitol in this system and those published earlier [1,2] in an NaF base solution. A similar direct comparison for sucrose [3,4] is more difficult because no capacity curves are available for the NaC1 base solution. Nevertheless Krishnan and de Levie [4] carried out a detailed study of the mutual effects of adsorption of sucrose and C1-.

Since the data for mannitol in NaF base solutions have been published in only very abbreviated form, it seems worthwhile to present these data in more detail and to a t tempt to account for the difference.

1Materials Development Division, AERE Harwell, Oxford OX11 0RA, United Kingdom.

0166-6622/90/$03.50 © 1990 --- Elsevier Science Publishers B.V.

50

EXPERIMENTAL

Differential capaci tance-potential curves were measured for eight concen- trations of mannitol in the range 0.031 mol dm -3 through to the saturated solution. The base electrolyte was 0.7953 mol dm -a aqueous NaF and the ref- erence electrode a 0.7953 mol dm -3 NaC1 calomel electrode. Electrocapillary curves were measured for seven of these mannitol concentrations, but the base electrolyte was 0.7953 mol dm -3 NaC1. Mannitol and NaF were B D H AnalaR grade and were used without further purification. NaC1 was thrice recrystal- lized from triply distilled water and, after drying in a vacuum desiccator, fused at red heat in a plat inum crucible. All the water used was triply distilled, the first stage being over alkaline KMnO4.

The differential capacitance was measured using methods described previ- ously [8,9]. The working electrode was a siliconized dropping mercury elec- trode of internal radius 0.04 mm. The mercury column height was typically 70 cm which produced an average mercury flow rate of 0.3 mg s - 1. The bridge was normally set to balance 6.8 s after the birth of a mercury drop whose total life was about 10 s. The measurements were recorded at a frequency of 800 Hz and a peak-to-peak amplitude of 10 mV. The capacity was independent of fre- quency over the range 0.4-6 kHz for all the concentrations studied. Electrode potentials were measured with a digital voltmeter accurate to + 0.5 mV. The electrolyte was freshly prepared and deoxygenated with presaturated oxygen- free nitrogen. The p.z.c, was then determined using the streaming electrode [10 ]. The capacity could not be measured at extreme negative potentials due to erratic behaviour of the dropping mercury electrode. This was characterised by the formation of a spray of t iny mercury droplets instead of a single spher- ical drop.

The electrocapillary curves were obtained using a high precision electrom- eter based on the maximum bubble pressure method [11,12 ]. This has been described in detail [5]. The capillary was unsiliconized. The instrument was calibrated with aqueous 0.1 M KCI solution, assuming a value for the interfa- cial tension at the electrocapillary maximum of 426.2 mN m-1 [13]. A more recent absolute determination by Vos and Los [ 14 ] shows that this value is too high by 0.6 mN m -1. However, this leads to a proportional error (0.15%) in the derived quantities which is quite negligible. The capillary constant was found to change systematically with use due to capillary wear and it was nec- essary to recalibrate the instrument frequently. Results in NaF base electrolyte were unreliable, especially at the more positive potentials, hence NaC1 base was used in this par t of the work. The densities of the working solutions were measured using a pycnometer that had been previously calibrated with triply distilled water.

RESULTS AND ANALYSIS

51

The differential capacity-potential curves are shown in Fig. 1. The shape is typical for the adsorption of an uncharged organic molecule with adsorption occurring in the region of the p.z.c, and a tendency towards desorption at the extremes of potential. However, it differs greatly from that observed for ali- phatic compounds [15] and is characteristic of a more repulsive interaction between the molecules in the adsorbed layer. This is also reflected in the rounded shape of the electrocapillary curves shown in Fig. 2 which were obtained with the electrometer. The conditions of these data will differ from those of the capacity data as a result of the specific adsorption of chloride ion in the former. In pure aqueous solution this has been well studied [16]. From these data, it may be estimated that chloride adsorption will be negligible in 0.8 M C1- so- lutions at potentials more negative than - 1.1 V versus the calomel electrode in the same solution (Fig. 3). If it may also be assumed that the presence of mannitol does not increase the adsorption of Cl- at the more negative poten- tials, then it is reasonable to use the experimental values of interfacial tension and charge at the most negative potentials in chloride solution as integration constants for the capacity data in fluoride solution. These were obtained from the data for the base solution. Integrations were performed using the program

3. ~

3C

/ Bo~e ?E , /3~,a ~M

56.6~M

,OOmM

25

/

2 C ~ " ' ~ /

I 0 0 .4 0 .8 1.2

- E / / V

Fig. 1. Differential capacity of a mercury electrode in contact with aqueous solutions of 0.7953 mol drn -3 NaF conta in ing manni to l at 25°C. The concentra t ion of manni to l is indicated by each curve.

52

430[

420

,ot/ Z ,

~- /

4 0 0

39C)ii B~SE IOO~M 178~M 316~M

380

-0-2 -0"4 -0"6 -0"8 -I-0 -1-2 m/v

Fig. 2. Electrocapillary curves for an aqueous solution of 0.7953 mol dm -3 NaC1 containing man- nitol at 25 ° C measured with the capillary electrometer. The concentration of mannitol is indicated by each curve.

-2o ! ilo'. 5

- 5

I -1-0

E/V

Fig. 3. Amount of chloride ion specifically adsorbed on Hg from aqueous KC1 at 25 °C (data from Ref. [16] ): ( × ) 1 mol dm-3; ( O ) 0.632 mol dm -3. The p.z.c, is indicated by vertical lines. Po- tential with respect to a 1 mol dm -3 KCI calomel electrode.

I I

53

TABLE 1

Coordinates of the p.z.c, of mercury in contact with 0.7953 mol dm -3 NaF containing mannitol at 25°C. Potentials are measured with respect to an aqueous 0.7953 mol dm -3 NaC1 calomel electrode

c /mmol dm -3 -Enz¢ /mV ypzc/mN m -1 Cpzc/l.tF cm -2

0 481 427.95 25.47 31.6 486 426.75 24.42 56.6 489 426.40 24.13

100 491 425.79 23.34 178 495 424.80 22.47 316 499 423.37 21.73 562 503 421.02 20.70

1014 509 417.74 19.75

described previously [17 ]. The resulting coordinates of the p.z.c, and the ca- pacity at this potential are collected in Table 1.

The maximum lowering of the capacity indicates a maximum adsorption at a small positive electrode charge value. The charge-potential curves intersect

M equal at a common point indicating that maximum adsorption occurs at amax to + 2.5 + 0.2/IC cm-2 (Em,x -- - 0.388 _+ 0.005 V ). The coincidence at this point was not as good as that for other sugar molecules; this may imply that the isotherm parameters describing the adsorption are dependent upon the elec- trical variable.

The surface pressure curves were congruent with respect to both electrical variables within the experimental error of the data. The composite isotherm, determined by superimposition of the data at other constant charges on to those at the charge of maximum adsorption, is shown in Fig. 4. The scatter of the data is of the order of + 0.5 mN m - 1 and is greatest for the lowest concen- trations. An error in the highest concentration may be indicated as all the values are consistently 0.5 mN m-1 high. The result shows that mannitol is only weakly adsorbed and is not approaching saturation. In order to study the adsorption more fully, it would be necessary to work at a higher temperature to increase the accessible concentration range.

Comparison of the logarithmic plot with the generalised isotherms shows that the virial, Volmer and Langmuir isotherms all fit the experimental curve with the parameters;

Volmer

log fl= 0.28

b -- 2.0.109 cm 2 mol -

Virial

54

logfl=0.78

B = 3.03-109 cm 2 mol- 1

Langmuir

logfl=0.41

Fs =3.34"10 -1° mol cm -2

The theoretical surface pressure curves calculated using these parameters and the corresponding integrated equations are shown in Fig. 4. However, the pa- rameter b for the Volmer equation has an unrealistic value as it corresponds to a molecular area on the surface of only 0.17 nm 2 which is inconsistent with the molecular models of Fig. 5 which shows possible configurations. The same is true for the virial coefficient which, in the absence of particle-particle inter- action, reduces to twice the molecular area and, therefore, corresponds to an area of 0.25 nm 2. The saturation coverage obtained for the Langmuir isotherm is consistent with the molecular model lying with the long axis flat on the electrode surface. However, the conformation of the molecule is unknown at the surface and so this conclusion can only be tentative. Although the Lang- muir isotherm parameters describe the surface pressure results reasonably well, they fail to reproduce the experimental surface excess data (see below).

The thermodynamic surface excesses were obtained by numerical differen-

I©

Z E

4

-2-5

,/

/7

-2.O -1.5 - I .o -0-5 O I_OGC + f(c~)

Fig. 4. Composite surface pressure curve at constant electrode charge for mannitol on a mercury electrode. Sets of points for other constant charges are superimposed on those for the charge of maximum adsorption by lateral displacement. Points are experimental; ( ) virial and Volmer equations; ( - - - ) Langmuir equation best fit; (- + - + ) Langmuir equation using parameters de- rived from the surface excess data.

55

{a) P L A N

H

dH2dH 2 12

HOCH I

3 13 SIDE VIEW,

HOCH 144

HCOH

HCOH

CH2©H

0 O-5nm s ide v i e w

(b) PLAN SIDE VIEW

41" side view 0 0-5

, . . . . . A M

Fig. 5. Mannitol space filling model: (a) with chain extended; (b) with chain coiled.

tiation of the surface pressure according to the Gibbs equation, using the pro- gram described previously [ 17 ] and assuming ideal behaviour for the adsorbate in the bulk. Recent experimental work on the ternary system sodium chloride-

56

3.0

2.0

I.©

©

TOI~M O0000OO

0 0 0 0 0 0 ~ * " ~ 0 °

- - ' - ' - - - ' " - . ' o "'

12 6 0 -6 -12

oM/1oCcm -~

Fig. 6. Surface excess o f m a n n i t o l as a f u n c t i o n o f charge on t h e electrode. Bu lk c o n c e n t r a t i o n is i nd i ca t ed by each curve. P o i n t s are der ived f rom e x p e r i m e n t u s i n g t h e Gibbs equa t ion . L i n e s are ca lcula ted as in Fig. 4.

manni tol -water shows that this assumption is probably justified [18]. The relative surface excess is shown as a function of the electrode charge and con- centration in Fig. 6. The theoretical curves are calculated using the parameters quoted above and the agreement is found to be poor. Maximum adsorption occurs at positive charge values and is consistent with the conclusions drawn from the charge-potential curves.

The Langmuir and virial isotherms were tested directly according to the linearized form:

c l Y = c l l - ' s + lll- 'sfl (1)

In (c /F) = 2 B F - In fl ( 2 )

respectively, and linear relationships were found in both cases. The isotherm parameters obtained from the slope and intercept of these graphs are different from those found from the surface pressure curves.

Langmuir

log f l= 0.14

/'s =4.43"10-1° mol cm -2

Virial

jc(a)

0

-0.1

-0 .2

- 0 . 3

-0 '4

- 0 " 5

- 0 " 6

57

-0 -7 10

- 0-7

-0-6

/ d ) -0 -5

-0 -4

-0"3

_0.2 I

-0-1

. . . . , . . . . , , , , , . . . . . L , 0

5 0 - 5 - 1 0

d M / p c c m -2

(b)

, , , • . . . . . . . . . . . . . .

0 6 0 1 2 0 1 8 0

( o M - dMmax)2/p C2c ~ 4

(c) (d) 0 - 0.7

- 0 , I -O.6

. { ( E ) S( E )

-0,2 -0"5

-0 '3 -0 .4

-0"4 -0-3

-0 .5 - 0 ' 2

-0 -6 -0.1

-0 .7 • • • i I i , i I 0 - 0 -5 - 110 0 0:1 0"2 0:3 0 .4

dv (E_EmoJ/V 2

Fig. 7. Variation of standard Gibbs energy of adsorption with electrical variable for mannitol adsorption: (a) electrode charge as independent variable; (b) expressed as quadratic function of charge: ( × ) charge more negative than a~x; ( O ) charge more positive than a~x; (c) potential as independent variable; (d) expressed as quadratic function: ( 0 ) potentials more negative than Era.x; ( X ) potentials more positive than Em~x.

58

logfl=0.34

B = 1.7-109 cm 2 mo1-1

The parameters for the virial equation are unrealistic as B, in the absence of interaction, corresponds to an area at saturation of 0.14 nm 2. The surface ex- cesses calculated for the Langmuir equation with the parameters recorded above show reasonable agreement with the experimental surface excess (Fig. 6). However, the agreement with the surface pressure is poor (Fig. 4). Therefore, the experimental results are not accurate enough to distinguish between the isotherms for this adsorbate. It should be noted that this becomes apparent only by comparing the experimental and theoretical calculation for both the surface pressure and the surface excess data. This demonstrates that an iso- therm should be tested in as many ways as possible before concluding that a particular model can describe the adsorption properties.

The uncertainty in the assignment of the adsorption isotherm does not allow a reliable value of the standard Gibbs energy of adsorption to be evaluated but, qualitatively, it is of the same order of magnitude as for xylose [5]. However, the variation of the standard Gibbs energy with the electrical variable is in-

(a)

- 0 . 5

>

~ o

0 ' 5 0

(b)

N 5 U

~)

~--o

- 5

- I O

a M / ~ c c n T 2

8

6

4

I i 1 0

I 1 2

1 .

I , I

0 -7 1.4 2. I io ' ° r /mot cm ~

2.8

E/;V C 1

0 6

G 7

0 0"7 I-4 2-1 2"8

lo'°r/mol crn ~

Fig. 8. (a) Shift of potential at constant charge as a function of amount of mannitol adsorbed. (b) Shift of charge at constant potential as a function of amount of mannitol adsorbed.

59

dependent of the isotherm and can be determined from the shift to superim- pose the surface pressure data onto the composite isotherm at constant elec- trical variable. The dependence is quadratic for both electrical variables and shows the same characteristic t rends as in the other systems (Fig. 7). A para- bolic fit (dotted line in Fig. 7a) is evident at constant charge when aM<o'Max and the constant b in the equation

log fl= log flma,, - b ( a--ama,,) 2 (3)

has a value of 0.0046 cm 4 ]~C-2. The dependence of the electrical variable on the amount adsorbed is linear in both cases (Fig. 8), confirming that the sys- tem is congruent, within the experimental accuracy, in either electrical vari- able.

DISCUSSION

A comparison with the data obtained on 1 mol d m - 3 NaC1 shows that despite the differences in the capacity-potential curves, the position of the maximum adsorption differs by only 0.5 ttC cm -2. This is particularly surprising since the specific adsorption of chloride ion in this solution must amount to over 12 ttC cm -2 in this region (see Fig. 3 ). Thus, a model in which the effective charge is regarded as the sum of the electrode charge and that on specifically adsorbed ions [19] could not be maintained in this system. This view is supported by the general similarity of the F(a) curves of Fig. 6 and those of Fig. 9 of Ref. [6]; the former are perhaps slightly more symmetrical than the latter, espe- cially at the lower concentrations. The displacement of C1- by mannitol seems unlikely because the amounts adsorbed at a given concentration differ rather little in the two base solutions, but only measurement of the surface excess of C1- in the presence of mannitol would confirm this. The saturation coverages estimated in the two systems differ more, but this may be a result of using different isotherms for their estimation.

The question of the mutual interaction of chloride ion and sucrose was ad- dressed by Krishnan and de Levie [4 ] in particular by measuring the amount of specifically adsorbed chloride in the presence of varying amounts of sucrose. The effect is small (a similar small effect was found by Silva when thiourea was adsorbed in chloride solution [20] ); it can also be seen that the charge at the adsorption maximum is even more closely similar for NaC1 and NaF base electrolytes than it is for mannitol. On the other hand, there are quite signifi- cant differences in the surface excesses of sucrose derived from these two data sets. At low surface excess of sucrose the values are greater in NaC1 than in NaF while at high surface excess they are less. In the latter region this differ- ence may be due to the assumption of ideality of the solutions made in Ref. [3 ] but it is unlikely tha t this would affect the former region. It also seems unlikely that computational errors could account for the differences here; the results

60

obtained in Ref. [4] must be more accurate in view of the larger number of concentrations studied.

It must, therefore, be concluded that the specific adsorption of halide ions has a small effect on the adsorption of mannitol and of sucrose. While this effect is quite significant for the capacity-potential curve of mannitol, it is not as large as would be expected from the simple electrostatic model of non-elec- trolyte adsorption originated by Frumkin [21] and Butler [22]. It is possible that the origin of this behaviour may lie in the fact that sugar molecules can replace water molecules in the water structure or in ionic solvation sheaths without any large energy changes. This is indicated by the nearly ideal behav- iour of these molecules in the ternary mixture [18]. It is clear that work on other related molecules and investigation by more molecular techniques such as infrared spectroscopy would be necessary to verify this hypothesis.

ACKNOWLEDGEMENTS

We are grateful to Drs R.M. Reeves and I. Williams who designed and built the maximum bubble pressure electrometer and to the Science Research Coun- cil for a maintenance grant to R. Peat.

REFERENCES

1 R. Parsons, in Z. Galus (Ed.), Adsorpcja na Elektrodach i Inhibitowanie Reakcji Electro- dowych, Warsaw, 1980, p. 7.

2 R. Parsons and R. Peat, J. Res. Inst. Catalysis, Hokkaido Univ., 28 (1980) 321. 3 R. Parsons and R. Peat, J. Electroanal. Chem., 122 (1981) 299. 4 M. Krishnan and R. de Levie, J. Electroanal. Chem., 131 (1982) 97. 5 R. Parsons and R. Peat, Proc. Indian Acad. Sci., 97 (1986) 333. 6 F. Silva and S. Rodrigues, J. Electroanal. Chem., 252 (1988) 403. 7 R. Peat and S. Shannon, J. Electroanal. Chem., 159 (1983) 229. 8 G.J. Hills and R. Payne, Trans. Faraday Soc., 61 (1965) 326. 9 R. Parsons, R. Peat and R.M. Reeves, J. Electroanal. Chem., 62 (1975) 151.

10 D.C. Grahame, E.M. Coffin, J.I. Cummings and M.A. Poth, J. Am. Chem. Soc., 74 (1952) 1207.

11 D.J. Schiffrin, J. Electroanal. Chem., 23 (1969) 168. 12 J. Lawrence and D.M. Mohilner, J. Electrochem. Soc., 118 (1971) 259; 1596. 13 M.A.V. Devanathan and P. Peries, Trans. Faraday Soc., 50 (1954) 1236. 14 H. Vos and J.M. Los, J. Colloid Interface Sci., 74 (1980) 360. 15 A.N. Frumkin and B.B. Damaskin, in J.O'M. Bockris and B.E. Conway (Eds), Modern As-

pects of Electrochemistry, Vol. 3, Butterworths, London, 1964, p. 149. 16 D.C. Grahame and R. Parsons, J. Am. Chem. Soc., 82 (1961) 1291. 17 J. Lawrence, R. Parsons and R. Payne, J. Electroanal. Chem., 16 (1968) 193. 18 R.V. Brill, T.R.C. Current Data News, 6 (1978). 19 R. Parsons and F.G.R. Zobel, Trans. Faraday Soc., 62 {1966) 3511. 20 F. Silva, Ph.D. Thesis, Southampton, 1981. 21 A.N. Frumkin, Z. Physik., 35 (1926) 792. 22 J.A.V. Butler, Proc. R. Soc. London, Ser. A, 122 (1929) 399.