Electroanalytical Chemistry and lnterfacial Electrochemistry, 60(1975) 183 196 (~ Elsevier Sequoia S.A., Lausanne - Printed in The Netherlands

183

SPECIFIC ADSORPTION OF SULFATE IONS AT A MERCURY ELEC- TRODE FROM AQUEOUS SODIUM SULFATE SOLUTIONS

RICHARD PAYNE

Air Force Cambridge Laboratories, Hanscom AFB, Bedford, Mass. 01731 (U.S.A.)

(Received 23rd December 1974)

INTRODUCTION

Specific adsorption of sulfate ions on mercury was investigated by Grahame and Soderberg 1 in a general study of adsorption from tenth-normal aqueous solutions. According to them the evidence for specific adsorption at the potential of zero charge (p.z.c.) is ambiguous. Thus they noted significant deviations from diffuse layer theory consistent with the occurrence of specific adsorption of anions, but found no corresponding negative shift of the p.z.c. On the other hand there is no doubt from these measurements that strong specific adsorption occurs on a positively charged electrode. In a recent paper Ivanov e t al. 2 have studied con- centrated solutions of ammonium sulfate, concluding that specific adsorption at the p.z.c, is absent and only minimal at a positively charged electrode. They did not however undertake a detailed analysis of their results.

In view of the ambiguities apparent in the previous work, and since no detailed account of the adsorption of a divalent anion has been given it seems appropriate to re-examine the behavior of the sulfate ion.

EXPERIMENTAL

The bridge and cell arrangement for the measurement of the double layer capacity at a dropping mercury electrode have been described previously a. The measurements were made at a single frequency of 1 kHz. The potentials were measured against a mercury-mercurous sulfate electrode in 0.1 M Na2SO 4 solution. The stability of the reference electrode was checked at frequent intervals using as a reference point the potential of zero charge for 0.1 M Na2SO 4. The p.z.c. was measured for each solution by the method of the streaming mercury electrode 4.

Ordinary distilled water was purified from trace organics by redistillation from alkaline permanganate. Analytical Reagent Na2SO4 was recrystallised, and the solutions prepared from this water by a volumetric method.

Capillaries were drawn from 0.5 mm bore pyrex capillary tubing. They were treated with dimethyldichlorosilane vapor as described previously 5 in order to prevent solution creep inside the capillary.

Oxygen was removed from the solution prior to admitting mercury into the cell by bubbling with oxygen-free nitrogen which had been pre-saturated with the solution. The cell and reference electrode were immersed in a water bath con- trolled to _+0.05°C.

1 8 4 R. P A Y N E

R E S U L T S

The double layer capacity is shown as a function of the potential in Fig. 1 for eight concentrations of Na2SO4 from 0.005 to 1 molar. The minimum due to the diffuse layer capacity evident in the lower concentrations occurs on the negative side of the p.z.c, as predicted by the Gouy~hapman theory for a 1:2 (NazSO4 type) electrolyte 6. The prominence of the diffuse layer capacity minimum suggests that specific adsorption of ions close to the p.z.c, is insignificant, at least for solutions more dilute than 0.1 M. Strong specific adsorption of anions is indicated however at more positive potentials by the sharp rise of the capacity. This is con- firmed by a detailed analysis of the data.

50

'~ 4C

l.,t.

8 3c a. o u

~ 2o C2~

I

%

o.o

I0 I "11

- 0.15 - I. 0 I. ~5 - 2 .0 Potential/V ( O.IM Hg/Hg2SO 4)

Fig. 1. Differential capac i ty curves for aqueous sod ium sulfate so lu t ions at the ind ica ted m o l a r con- cen t ra t ions and 25~'C.

Molar concentrations were converted to molalities using density data taken from the International Critical Tables. Activity coefficients were then interpolated from published values. The measured potential values were converted to a scale (E-) based on a sulfate ion reversible reference electrode by subtracting an amount ( 3 R T / 2 F ) t + ln(:~ ±/~o.1 M), the e.m.f, of the cell with transport

Hg,Hg2SO4 1 NazSO4 ,I Na2SO4 I HgzSO4, Hg 0.1M i x M

A concentration independent Value of 0.3802 was taken for t +, the cation transport number 7. The capacity was integrated twice on a computer (Elliott 803 or CDC6600) using the measured potential of zero charge as the first integration

A D S O R P T I O N OF SO4 z ON Hg IN H 2 0 185

constant. The second integration constant was determined from Gouy's interfacial tension data for NazSO ~ and other sulfate solutions s. The capacity, potential and interfacial tension (7) were interpolated at integral values of q, the electrode charge, by the computer program. The surface excess of cations F+ was obtained, for positive values of q, by graphical differentiation of plotted curves of 7 + q E - against chemical potential of the electrolyte at constant q. For negative values of q, the surface excess of anions was determined using the analogous plot of 7 + qE + against chemical potential.

The coordinates of the electrocapillary maximum are given in Table 1.

TABLE 1

C O O R D I N A T E S OF THE E L E C T R O C A P I L L A R Y M A X I M U M FOR THE M E R C U R Y - A Q U E O U S NazSO, ~ INTERFACE AT 25°C

c / m o l l - 1 - - E ~ / V 7~/erg c m - 2 b C Z / p F c m - 2

raeasd, a vs. N C E

0.005 0.876 0.474 425.3 14.16 0.01 0.870 0.475 425.4 16.47 0.02 0.863 0.476 425.4 18.79 0.05 0.856 0.479 425.5 21.61 0.1 0.852 0.481 425.6 23.50 0.2 0.848 0.483 425.7 25.02 0.5 0.843 0.486 426.1 26.85 1.0 0.842 0.490 426.4 28.20

a Experimentally measured value includes liquid junct ion potential between the solution and a 0.1 M Na2SO 4 solution. b Interpolated from Gouy's data s for NazSO4 solutions at 18°C corrected to 25°C.

DISCUSSION

The minimum in the diffuse layer capacity The diffuse layer theory for a 1 : 2 electrolyte leads 6 to the following equations

for the charge (qd) and the differential capacity (C d) of the diffuse part of the double layer:

qa = _A(e r _ l ) ( l + 2 e Y) (1)

C d = ( A e ° / k T ) ( e 2 r - e - r ) / ( e r - 1)(1 + 2e-Y)~ (2)

where e ° is the electronic charge and k the Boltzmann constant. The concentration dependent parameter A is given by

A = (e kTni/27z) ~ (3)

where e is the dielectric constant of the solvent and ni is the concentration of the less abundant ion (i.e. SO 2-) in ions per cubic centimeter of solution. The remaining parameter y is related to the potential difference across the diffuse layer q52 according to

y = e ° 4)2/kW (4)

At 25°C y= 38.92~b 2.

186 R. PAYNE

40C

300

Cd /~1 IV-'

200

I00

-,o -~ 6 ~ ,~ qdlA

Fig. 2. Reduced differential capacity of the diffuse la)er as a function of the reduced diffuse layer charge according to the G o u y q S h a p m a n theory for a 1:2 electrolyte at 25C .

Unlike the case of the symmetrical electrolyte, C a in this case cannot be expressed as an explicit function of qd. However the equations can be solved numerically with the result shown in Fig. 2. As Grahame pointed out, the minimum in the diffuse layer capacity for this type of electrolyte occurs on the negative side of the p.z.c. The coordinates of the minimum in Fig. 2 are qd/A= 1.021, Ca/A = 60.87 V- 1 and ~b2 = -0.0162 V. It is interesting to note that while the position of the capacity minimum shifts to more negative values of the diffuse layer charge with increasing concentration, the value of ~b 2 at the minimum is independent of the concentration.

The experimental capacity curves in Fig. 1 show clearly that the diffuse layer minimum occurs on the negative side of the p.z.c, as predicted by the theory. Experiment and theory were compared quantitatively for the most dilute solutions by integrating the capacity from the p.z.c, to the minimum. The minima occur at values of qd/A of 1.107, 1.104 and 1.213, respectively, for the three most dilute solutions in increasing order of concentration, compared with the theoretical value of 1.021. The observed discrepancy could result from the displacement of the minimum by the effect of the steeply rising inner layer capacity superimposed on the diffuse layer capacity. Therefore the actual diffuse layer capacity was calculated from the experimental results with the aid of Grahame's series capacitor formula 9 for the total capacity C:

1/C = 1/C' + 1/C d (5)

assuming the absence of specific adsorption, where C i is the inner layer capacity. Values of C ~ were calculated from the results for the 0.1 M solution. The diffuse layer capacity was then calculated from the experimental results for the dilute solutions using eqn. (5). The results shown in Fig. 3 confirm the initial conclusion. The minimum in the capacity apparently occurs at a more negative value of the electrode charge than predicted by the theory. However the capacities are also lower than the theoretical values, the discrepancy increasing with the concentration.

ADSORPTION OF SO~ ON Hg IN H20 187

I i

31

3O

'~9 29

28

27

26

25 t • i

I I

0 -0.5 q/}JC cr6 2-LO

Fig. 3. Comparison of experimentally determined diffuse layer capacities (circles) with values calculated (lines) from the diffuse layer theory for two concentrations of Na2SO~. The method of calculation is indicated in the text.

The divergence of both the position of the capacity minimum and the magni- tude of the capacity from the theory can be explained in terms of specific adsorption of anions which so far has been neglected. However another factor which may be quite important is ion-pairing which would result in lower values of C d. In view of the uncertainties introduced by the various assumptions made in the analysis, the agreement between theory and experiment must be considered satisfactory.

Specific adsorption of sulfate ions at the p.z.c. The evidence for or against specific adsorption of sulfate ions on an un-

charged mercury electrode is somewhat ambiguous. For example Grahame and Soderberg 1 noted significant deviations from the diffuse layer for a nonadsorbed electrolyte in the case of 0.05 M K z S O 4 but found no shift of the p.z.c. The results of Gouy 8 for a number of sulfates and the recent measurements of Ivanov et al. for ( N H 4 ) 2 S O 4 solutions show clearly that surface deficiencies of electrolyte occur at the p.z.c. Ivanov et al. conclude that the elevation of the interfacial tension at the air~olution and the uncharged mercury~olution interfaces is essentially the same. However their measurements seem to show a much larger effect at the air-solution interface, which suggests that some specific adsorption may be occurring at the electrode.

The variation of the interfacial tension with electrolyte concentration can be calculated for a model of total exclusion of electrolyte from the first layer of solvent with the aid of certain assumptions. In the simplest case the density of the water in the inner layer is assumed equal to the bulk value. The number (nw) of water molecules in unit area of a monolayer is then given by

188 R. PAYNE

, , , = ( U / V ? (6) where N is Avogadro's number (5.023 × l023) and V is the molar volume of water (18 cm3). The value of nw therefore is 1.040x l01~ molecule cm -2. The thickness of the monolayer (x2) is given by

x2 = ( V / N ) ~ (7)

so that x 2 is equal to 3.10 × 10 s cm. The electrolyte associated with n~ molecules of solvent in the bulk solution is the amount (E,,~,) excluded from the monolayer and is given by

Frs,l , = ( z M e ° / / 5 5 . 5 5 ) ( N / V ) z3 (8)

= z M" 3.00 I~C cm 2

where M is the molality of the electrolyte and z is the valence of the less abundant ion (i.e. SO4 z ).

The predicted variation of the interfacial tension (AT) was calculated from the integral form of the Gibbs adsorption equation,

t l n a ±

A7 = 3 R T La,t d In a _ (9)

by numerical integration using E,,I, values given by eqn. (8). The resultant curve is compared in Fig. 4 with the experimental data of Gouy, and Ivanov et al. Gouy's data points for Na2SO 4 solutions evidently lie on the curve within the experimental error. The mercury-(NH4)2SO4 data of Ivanov et al. on the other hand appear to lie well below the curve at the higher concentrations*. Their data for the air-solution interface are in much better agreement with the calculated values as are the corre- sponding results of Jones and Ray 1° for K2SO 4 solutions. The experimental results therefore appear to be broadly consistent with the presence of an electrolyte-free monolayer of water and provide no clear evidence for specific adsorption of anions.

As a further test for specific adsorption the concentration dependence of the capacity can be compared with the predictions of the diffuse layer theory. This can be done most conveniently by plotting 1/C against 1/C d at constant q according to eqn. (5). In the absence of specific adsorption the plot should be linear with a slope of unity. The results of this plot shown in Fig. 5 indicate some deviation from the line of unit slope at the higher concentrations for q >~ 0 consistent with specific adsorption of anions. This test of course depends on the validity of the series capacitor model and the diffuse layer theory which is questionable at the higher concen- trations of the electrolyte.

More reliable evidence for specific adsorption of anions however is provided by the variation of the potential of zero charge with concentration as shown in Fig. 6. The p.z.c., referred to the normal calomel scale, shifts strongly in the negative direction with increasing concentration. At low concentrations the p.z.c, approaches a limiting value close to -0 .472 V, the established value for a nonadsorbed electrolyte la. The value -0 .4705 V for a 0.05 M KzSO 4 solution reported by Grahame et al. ~'~2 appears to be too positive by ~8 inV. Included in this value

* However the value of 3 × 10 -s cm obtained by Ivanov et al. for the thickness of the electrolyte- free layer is close to the value of 3.10 × 10 -s calculated from eqn. (7).

ADSORPTION OF SO~ ON Hg IN H20 189

2

<~

0

-0120

NazSO 4 - GouyS

(NH4)zSO 4 -I vonov et ol2

,, ,, - air-solution z

0.09i

D.07

( -I/

c m 2 pF"

0.05

3.0~

q=O

-8

' I ' ' ' ' b ' -0.15 -010 -0.0. ~ 0 0.02 O. 4 0.06 {3RT/F)InmT_+ / V (cd)-I/cm2 ,uF-I

Fig. 4. Variation of the interfacial tension with electrolyte concentration. Points are experimental values. The line is calculated from eqn. (9) assuming a value of 425.4 erg cm- 2 for the limiting value of the interfacial tension at infinite dilution.

Fig. 5. Reciprocal of the measured capacity plotted against the reciprocal diffuse layer capacity according to eqn. (5).

500

49~

-Ez/ rnV

48C

4?C

I i , i

0 This work

• Ivonov et ol2- (NH4)2S04 • v Grahame etal.q I z - K2SO 4 •

• K2SO 4 - n o LdPcorrn.

x Grahams 9 - N ~ ~ /

, ? -d2o -o,5 -olb -o6~

(SRTIF)InmT± / V

Fig. 6. Potential of zero charge on the normal calomel scale as a function of the concentration for NazSO 4 and (NH4)zSO4 solutions.

however is a correction of 6.1 mV for the liquid junction potential between 0.05 M K2SO 4 and 0.1 M KC1 which introduces some uncertainty. The present results and those of Ivanov e t al. 2 were obtained with a sulfate reversible electrode and for that reason are more reliable. The agreement between the NazSO 4 and

190 R. PA3NE

(NH4)2SO 4 data is excellent up to -0 .2 M. At higher concentrations the (NH4)2SO4 potentials are more negative. The difference may be due to different degrees of ion-pairing. For example 0.1 M solutions contain approximately 10°o of NaSO4 and 20°/o ofNH4SO 2 ~ 3. At high concentrations ion-pairing in (NH4)2SO 4 solutions is evidently so extensive that the solution behaves effectively as a 1:1 electrolyte 14. This factor could readily explain the results m Fig. 6. It is clear from Gouy's measurements 8 that the H S O 4 ion is quite strongly adsorbed. It seems reasonable to assume therefore that the NH4SO,~ ion would also be adsorbed.

Grahame and Soderberg ~ developed a further piece of evidence pointing toward specific adsorption of anions at the p.z.c. They determined a transference number ~+, the fraction of the charge carried into or out of the double layer by cations, which they compared with the value predicted by the diffuse layer theory for a nonadsorbed electrolyte. However this is not an independent test since T + is related to the potential of zero charge. The transference number r ~ (and the corresponding parameter for anions, r ) is actually a form of the Esin and Markov coefficient. For a 1:2 type electrolyte the relationships are

and,

where diffuse

and,

"c+ = (dE-//?/0 q = -(,~r+/~qL (10)

T_ = (~E+/~B)q = -(i~F /?q),, (11)

/, is the chemical potential of the salt. For a nonadsorbed electrolyte the layer theory leads to,

-F (cF+/cq ) , = e - " / ( l + e " + e ") (12)

- e ( e F _ / g q ) , = (1 +eY)/(1 + e " + e - " ) (13)

At the p.z.c, therefore, y = 0 and a plot of E - against (3R T 2 F ) In a :~ should be linear with a slope of 0.333 equal to r , . The plot of the experimental data in Fi N 7 however is nonlinear and the slope is systematically lower than the theoretical value, consistent with the presence of specifically adsorbed anions. The slope at the 0.05 M concentration is 0.23. Grahame and Soderberg I reported a value of 0.261 for 0.05 M K2SO4, concluding that the deviation from diffuse layer behavior was signi- ficant. However this is not consistent with their finding that the p.z.c, is not shifted, which as mentioned earlier is probably due to the uncertainty of liquid junction potentials in their experimental arrangement. This inconsistency was ap- parently overlooked by them.

Specific adsorption of su!fate ions on a positirely charged electrode According to Ivanov et al. 2 specific adsorption from (NH4)2SO 4 solutions is

minimal even at the limit of positive polarisation. On the other hand Grahame and Soderberg a showed that specific adsorption from a 0.05 M KzSO ~ solution is strong, comparable for example to the adsorption of Cl- although shifted to somewhat more positive values of the electrode charge. The occurrence of strong specific adsorption of anions in the present system is confirmed by the plots of F+ against q which show the familiar minimum even at the lowest concentrations (Fig. 8). These curves are related to the potential dependence of the amount adsorbed which is exceptionally

ADSORPTION OF SO~- ON Hg IN H20 191

3£

l 2(

FF÷, - 0 . 8 0 ~uC e d

E-/V/ I0

-0.85 0 0 u

-0.90

0.00

I

i I I -o.zo -ols -o.Io -o.(~s -,e 20 I0 0 -I0 (3RT/F)Inm)'_+ IV q /)JC cni 2

Fig. 7. Esin and Markov plot for q=0. Points are experimental. Straight line of slope 1/3 is obtained from the diffuse layer according to eqns. (10) and (12). Slope given by the experimental points at the 0.05 M concentration is 0.23.

Fig. 8. Surface excess of cations as a function of the electrode charge for Na2SO4 solutions at the molar concentrations shown.

strong, contrasting sharply with the behavior of other polyatomic anions such as NO 3 15,16, C102 iv, and PF 6 18.19 in aqueous solutions.

An indication of the extent of the specific adsorption of SO 2- ions is provided by the Esin and Markov coefficient, (cE+/c,lt)q. The deviation of this coefficient from the diffuse layer theoretical value for a nonadsorbed electrolyte has already been noted for the special case of q=0. The plots of E + against # are shown for other values of q in Fig. 9. These curves are unusual for systems of this type inasmuch as they are markedly curved except for the extreme negative values of q. The slope is equal to the Esin and Markov coefficient, which according to the diffuse layer theory for a nonadsorbed electrolyte (eqns. 12 and 13) approaches the limit of zero at negative values of q (y large and positive). For the NazSO4 system in contrast the Esin and Markov coefficient reaches values more than twicethe diffuse layer limiting value on the positive side as shown in Fig. 10. This compares with reported values for other aqueous electrolytes of -1 .00 for KNO316, -1 .36 for KI 2° and - 1.22 for KC12°. The explanation of the unusually large effect in NazSO 4 is unclear. This factor is probably related to the sharp increase of the capacity on the anodic side of the p.z.c.

The specifically adsorbed charge was calculated with the aid of the diffuse layer theory assuming that cations are not specifically adsorbed on the positive side of the p.z.c. Since the exclusion of electrolyte from the inner layer is clearly an important factor at concentrations higher than ~ 0.1 M a correction is necessary in order to obtain realistic adsorption isotherms. In order to correct for the electrolyte exclusion effect an amount equal to 3.00 zM/~C cm-2 given by eqn. (8) was added

192 R. PAYNE

-0.5

E+/ V

-I.0

-I.5

-2£

o ~ - 4

-0" 0 0 0 ~ 0 " - - - 0 ' ~ -8

-0 0 0 0 ~ 0 - ~ 0 - - -12

0 0 0 ~ -IE

- .2o -6.,s -&o -8.o (&RT/2F) Ina_+ / V

Fig. 9. Esin and Markov plots for Na2SOa solutions. The electrode charge q (#C c m 2) is indicated for each line.

r i

I

20

1.5 0.1 *

FaE +

0 . 0 0 5

1.0

0.5

0 20 I0 0 - I0

qm /,uC cnT 2

Fig. I0. The Esin and Markov coefficient as a function of the electrode charge for NazSO4 solutions at the indicated molar concentrations.

to the thermodynamic surface excess of cations. A similar correction was applied to the anion surface excess. The corrected values of F+ were then used to calculate the diffuse layer concentration of anions and hence by difference the specifically adsorbed

ADSORPTION OF S O l - ON Hg IN H20 193

anionic charge. Implicit in this procedure is the assumption that eqn. (8), which seems to predict the variation of the interfacial tension at the p.z.c, is also valid for other values of q. This is obviously an uncertain extrapolation. It also ignores the fact that the surface concentration of water decreases as the water molecules are replaced by specifically adsorbed ions. Thus at saturation coverage of the electrode by specifically adsorbed ions the correction for electrolyte exclusion vanishes. The effects of electrostriction of solvent molecules by specifically adsorbed ions and ion-pairing further complicate matters. Ion-pairing in the diffuse layer will have the effect of decreasing the diffuse layer corrections to the surface excess and lowering the specifically adsorbed charge.

u =k

o"

2c

o. 10

i

O' // • .0" I,

t0.

- 0.20

i i q=20 [

.0

;0

-q

pC c i z

-0.15 -OlO -0.05

(~3RT/F)In my+ / V

i i i I

0.5

0.2 0.~.1

0

I I

0 I0 20 q /)JC crn -2

Fig. 11. Adsorption isotherms for specifically adsorbed sulfate ions at constant electrode charge. Filled circles are values corrected for electrolyte exclusion effects as described in the text. Open circles are uncorrected values.

Fig. 12. Specifically adsorbed charge due to sulfate ions as a function of the electrode charge at the molar concentrations shown.

The adsorption isotherms calculated by this rather crude approach are shown in Fig. 11. For comparison the results obtained without correction for salt exclusion are also shown. For q < 10 #C cm-2 neglect of this factor results in negative isotherm slopes which have no physical significance. As frequently found for inorganic anions the specifically adsorbed charge becomes asymptotic to the value of the electrode charge at low electrolyte concentrations. The plots of specifically adsorbed charge against q are shown for several concentrations in Fig. 12. These plots are remarkable only for the unusually large slope of about - 1.9 in the most concentrated solutions. This is related to the large Esin and Markov coefficient remarked upon earlier.

The plots of potential difference across the inner layer (~b"- 2) against speci- fically adsorbed charge (ql) at constant q were found to be linear and parallel for

194 R. PAYNE

values of q > 14 /tC cm-2. The partial differential capacities of the inner layer (Cql/~om-Z)q a n d ( c q / c ~ b n ' - 2 ) 1 in this range are approximately 160 I~F cm 2 and 25/LF c m z respectively, leading to a value for the thickness ratio (Xz-X~) 'x2 of 0.16. The results for q < 12 are unreliable due to the uncertainty of the diffuse layer potential difference.

Deviations from dti{'lilse layer behat'ior on a negatit:ely charged electrode In a previous paper 16 systematic deviations of the ionic surface excesses from

diffuse layer theory were noted for aqueous solutions of KNO3, NaOH, and NazSO 4 at negative values of the electrode charge. The effect is illustrated b) the Na2SO4 results shown in Fig. 13. The anion surface excess does not reach the limiting repulsion value predicted by diffuse layer theory but instead passes through a minimum.

I 0

5

FI-'_ /

.#C c ~

I

o O

o9__~_~_~ ~ - o ,

- L i ~ w I © 0 0 5

0 - 5 - I0 -15

q / p C crff 2

Fig. 13. Surface excess of anions on a negativel) charged electrode for three concentrations of the electrolyte. Lines are calculated from the diffuse layer theory assuming the absence of specific adsorption. Points are experimental.

increasing again at the extreme of negative polarisation. The appearance of these curves suggests (by analogy with the F . curves in Fig. 8) that the effect is due to specific adsorption of cations, although as noted elsewhere2122 the concentration dependence does not seem to be consistent with this explanation. The effect actually appears to be more pronounced in dilute solutions, especially in the case of NaOH. Parsons and Trasatti 23 suggested therefore that the effect might be due to slow relaxation of the diffuse layer resulting in low values of the measured capacity :4. However they were unable to confirm this experimentally. Furthermore similar effects are present in some results based on electrocapillary measurements only 25 where nonequilibrium effects could not be a factor.

A close examination of the literature shows that this effect occurs commonly in both aqueous and nonaqueous solutions. For example Frumkin et al. 26 found systematic deviations from diffuse layer behavior for 1 M aqueous solutions of KC1

ADSORPTION OF SO4 z- ON Hg IN H20 195

and CsC1 which they attributed to specific adsorption of cations, quite strong in the case of Cs + but weak in the case of K +. Similar results for 0.1 M solutions of NaC1, KC1 and CsCI in formamide 27 were also attributed to specific adsorption of cations. However deviations from diffuse layer behavior found for a 0.1 M solution of CsC1 in N-methylformamide 2s were attributed to dielectric saturation effects. More recently Parsons and Trasatti 23 found a substantial effect for dilute chloride solu- tions in water, while Schiffrin 29 noted a smaller but definite effect for 0.1 M and 0.1 M solutions of KF in water at 0°C and 15°C. Well-defined minima in the curves of F against q were found for concentrated solutions of KF 3° and CsF 31 and also (NH4)2SO~ 2. In the case of CsF the results were attributed to specific adsorption of Cs +, which however it was noted did not produce a sharp increase of the capacity.

There now seems little doubt that these results provide authentic evidence for the specific adsorption of cations for which there is abundant indirect evidence from electrode kinetics 32. The ambiguous concentration dependence can be explained by the weakness of the adsorption and the uncertainty introduced by electrolyte ex- clusion effects. Thus one may conclude that the sharp distinction drawn by Grahame 9 between specifically and nonspecifically adsorbed ions does not exist, and that all ions exhibit some degree of specific adsorption under sufficiently favorable conditions.

CONCLUSIONS

I. The minimum in the diffuse layer capacity in dilute Na2SO# solutions is located on the negative side of the p.z.c, as predicted by the Gouy~2hapman theory for a 1 : 2 electrolyte. The magnitude of the capacity at the minimum and the corresponding value of the ionic charge in the diffuse layer are in satisfactory agreement with the theory.

2. Specific adsorption of SO~- ions occurs at the p.z.c, as shown by the dependence of the p.z.c, and the related Esin and Markov coefficient on the elec- trolyte concentration. The interfacial tension however provides no evidence for speci- fic adsorption at the p.z.c, and is consistent with the presence of an electrolyte-free monolayer of water on the electrode.

3. Strong specific adsorption of SO 2- ions occurs on a positively charged electrode at all concentrations. The adsorption is characterised by an exceptionally large Esin and Markov effect.

4. Deviations from diffuse layer behavior observed for a negatively charged electrode confirm and extend similar results for various other systems reported in the literature and are consistent with specific adsorption of cations.

ACKNOWLEDGMENT

The experimental part of this work was performed at the University of Bristol during the tenure of a Senior Fellowship of the Science Research Council.

REFERENCES

1 D. C. Grahame and B. Soderberg, J. Chem. Phys., 22 (1954) 449. 2 V. F. Ivanov, B. B. Damaskin, B. S. Segelman and N. I. Melekhova, Elektrokhimiya, 9 (1973) 389.

196 R. PAYNE

3 G. J. Hills and R. Payne, Trans. Faraday Sot.. 61 (1965) 316. 4 D. C. Grahame, R. P. Larsen and M. A. Poth. ,I..4mer. Chem. Soc.. 71 (1949) 2978. 5 R. Payne, J. ElectroanaL Chem., 7 (1964) 343. 6 D. C. Grahame, J. Chem. Phys., 21 (1953) 1054. 7 R. Parsons, Handbook ~[ Electrochemical Constants. Buttervvorths. London, 1959. 8 G. Gouy, AmT, Chim. Phys.. 29 (7)(1903) 145. 9 D. C. Grahame, Chem. Ret:.. 41 (1947) 441.

10 G. Jones and W. A. Ray, J. Amer. Chem. Sot., 59 (1937) 187. 11 D. C. Grahame, J. Amer. Chem. Soc., 76 (1954) 4819. 12 D. C. Grahame, E. M. Coffin, J. I. Cummings and M. A. Poth. J. Amer. Chem. Soc.. 74 (1952) 1207. 13 C. W. Davies, Ion Association. Butterworths, Washington. 1962. p. 42. 14 R. A. Robinson and R. H. Stokes. Electrolyte Salution.s. Butterworths. London. 1955. p. 372. 15 R. Payne, J. Phys. Chem., 69 (1965)4113. 16 R. Payne, J. Electrochem. Soc.. 113 (1966) 999. 17 R. Payne, J. Phys. Chem.. 70 (1966) 204. 18 G. J. Hills and R. M. Reeves. J. Electroanal. Chem.. 31 (1971) 269. 19 L. M. Baugh and R. Parsons, J. Electroanal. Chem.. 40 (1972)407. 20 R. Parsons, Proc. 2nd International Congress o[' SmJ~lce Actility, London. 3 (1957) 38. 21 R. Payne, J. Eleetroanal. Chem., 41 (1973) 277. 22 R. Payne in J. F. Danielli, M. D. Rosenberg and D. A. Cadenhead (Eds.). Progress m Sur(ace and

Membrane Science, Vol. 6, Academic Press, Ne~ York. 1973. p. 51. 23 R. Parsons and S. Trasatti. Trans. Faraday Soc.. 65 (1969) 3314. 24 G. C. Barker, J. Electroanal. Chem.. 12 (1966) 495. 25 L. F. Oldfield, Ph.D. Thesis, London. 1951. 26 A. N. Frumkin, R. V. Ivanova and B. B. Damaskin. Ookl. Akad. Nauk SSSR. 157 (1964) 1202. 27 B. B. Damaskin, R. V. Ivanova and A. A. Survila, Elektrokhimiya, 1 (1965) 767. 28 B. B. Damaskin and R. V. lvanova, Zh. Fiz. Khim.. 38 (1964) 176. 29 D. J. Schiffrin, Trans. Faraday Soe., 67 (1971) 3318. 30 N. I. Melekhova, V. F. Ivanov and B. B. Damaskin. Elektrokhimiya, 5 (1969) 613. 31 B. B. Damaskin, V. F. Ivanov and N. I. Melekhova. Elektrokhimiya, 6 (1970) 385. 32 R. Parsons in P. Delahay and C. W. Tobias (Eds.), Advances in Electrochemistr 3 amt Electrochemical

Engineering, Val. 1, Interscience, New York, 1961. p. 1.