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Multi layer adsorption at mercury electrodes
J. Electroanal. Chem., 66 (1975) 77--80
77
© Elsevier Sequoia S.A., Lausanne -- Printed in The Netherlands
Preliminary note MULTI LAYER ADSORPTION AT MERCURY ELECTRODES
P. G R U N D L E R and H. WENDLANDT
Karl-Marx-UniversitEt Leipzig, Sektion Chemie, 701 Leipzig, Liebigstr. 18 (D.D.R.) (Received 28th August 1975)
Adsorption processes play an important role in many fields. For practical applications, the surface concentration (surface excess) F of adsorbed species at metal/solution interfaces is of interest. In the course of investigation of some benzenethiosulphonate derivatives a very unusual behaviour of the substance p-acetamino-benzenethiosulphonate (in following shortened to ABTS) was found. NH.CO.CH 3
S02S-
Adsorption of this substance at mercury electrodes was studied by chronocoulometry. Some measurements by a.c.-polarography and galvanostatic pulses were also performed. ABTS is able to form poorly dissociated mercury salts. For substances of this type, Anson and co-workers [1--3] introduced chronocoulometry as a tool for investigation of adsorption processes. Measurements are performed by impressing three potentiostatic voltage steps to the electrode. Primarily, the electrode is held at a potential El (-1.0 V vs. SCE). At this potential, as was indicated by a.c.-polarography, neither adsorption nor any electrochemical reaction occurred. In the first step, the fixed pbtential El is instantaneously switched to a variable potential E2, which permits adsorption, but no electrochemical reaction. Integration of the current flowing during this step gives the charge Q1, which contains double layer charging as well as partial charge transfer (according to Lorenz [4] ). In a second experiment E2 is switched to a fixed potential E3 (+0.2 V vs. SCE). In this step anodic formation of mercury salts occurs. This was proved by d.c.-polarographic studies which established the existence of an anodic polarographic wave with a half wave potential of +0.15 V vs. SCE. Electrolysis currents during this step are integrated and plotted against the square root of
78 time. By extrapolating the straight line of this plot to t = 0 the charge Q2 is obtained. In the same w a y a charge Q3 was determined from a third experiment during which E~ was stepped to E3 directly. From eqn. (1) the a m o u n t of reactant r (mol cm -2) adsorbed on the electrode surface at potential E2 can be calculated, if n (number of electrons involved in anodic salt formation) is known. n F F = Q I (EI-~E2) + Q2 (E2 ~E3 ) - Q3 ( E l s E 3 )
(1)
The value of n was determined b y a coulometric experiment, during which the potential E3 was impressed to a large mercury pool electrode in stirred solution. From the integrated electrolysis current and reactant concentration decrease in the course of electrolysis n was found to be equal 1, i.e. any molecule ABTS was coupled to a one-electron change in anodic salt formation. In our chronocoulometric experiments a simple electronic circuit was used which consisted mainly of an operational amplifier type 741 connected as a potentiostat. An ordinary switch coupled with a small damping capacitor was connected to the non inverting input and allowed to switch in arbitrarily adjustable reference voltages thus producing single pulses at the o u t p u t of the amplifier. The inverting input was connected to the reference electrode (SCE). The o u t p u t of the amplifier was connected through an integrating capacitor to an auxiliary electrode in the cell. Voltage changes of the integrating capacitor were followed b y an oscilloscope, or a recorder in the case of measuring QI. Influence of the offset of the amplifier current was graphically eliminated. Overall response time of the electronic arrangement was 75 ~s. The working electrode was a Kemula type mercury drop electrode with an area of 0.0362 cm 2 . The reference electrode was separated from the cell b y a salt bridge filled with 1 M KNO3. All solutions used were 1 M in KNO3. They also contained a borate (for pH 6.6) or a phosphate (for pH 5.0) buffer system. Values of the adsorption potential E2 varied from - 0 . 8 V to - 0 . 1 V vs. SCE. Concentrations of ABTS lay b e t w e e n 10 -3 M and 10 -2 M. Some chronocoulometric results are shown in Fig. 1. As can be seen from this figure, at potentials more negative than - 6 0 0 mV essentially no adsorpt-ion takes place. From this point to more positive potentials, the adsorbed a m o u n t shows a sharp increase which is followed by a nearly constant region. This is in agreement with a.c. polarograms, which show increasing differential capacities beginning at the same point. In the constant region F depends only on the solution concentration. Further increase in the measured F values in the direction to positive potentials may be due to the beginning of anodic salt formation during m e a s u r e m e n t of QI. The d.c. polarograms indicate slightly rising anodic currents beginning with - 0 . 1 V (for the lowest concentration of ABTS) and shifting to - 0 . 3 V (for the highest reactant concentration). These anodic currents tend to give high results in F. From these measurements as well as from a.c.-polarographic investigations and chronocoulometric studies of other benzenethiosulphonates which do n o t
79
v v
,
/
~' 3C E u E 9 ~ 2O
10
4
f AooJ
/'f f o ~
-10
- O15
° ~ o ~o . ~ r - ~ °
E2//V
vs. SCE
-0
Fig. 1. Surface concentrations of p-acetaminobenzenethiosulphonate solutions. Electrolyte: 1 M K N O s , 0.1 M borate buffer, p H = 6.6. (~) 1 x l0 -3 M, (~) 2 x 10 -3 M, (o) 4 x 10 -3 M, (V) 1 X I 0 -2 M .
contain an acetamino group [6] it is clear, t h a t anions of such compounds are strongly chemisorbed at mercury when potentials are less negative t h a n - 0 . 6 V. In a sharp contrast to the above mentioned studies of other benzenethiosulphonates are the very high values of F f o u n d with ABTS. In the case of benzenethiosulphonate, for instance, F reaches m a x i m u m values of n o t more than 5 x 10 -1° mol cm -2 . This corresponds to a surface area of a b o u t 35 A 2 occupied by each molecule. These data indicate a completely covered electrode surface. As can be seen from Fig. 1, in the case of ABTS, the electrode surface must be completely covered at potentials more positive than - 0 . 6 V already with the lowest reactant concentration studied. This statement holds even if an arrangement of the molecules perpendicular to electrode surface is assumed. For verification we calculated the area covered by one molecule. At reactant concentrations greater than 2 x 10 -3 M, more than one layer of molecules must exist at the electrode surface. We believe t h a t this occurs by a multiple layer adsorption mechanism, for which we propose the following explanation. The thiosulphonate group of the molecule is strongly acidic. Consequently nearly all molecules exist in the form of anions in neutral or slightly acid solution. Anion groups are subject to strong adsorption. On the other hand, any molecule is able to form cations with its acetamino group. One may assume, therefore, that adsorbed molecules form a layer of "cationic molecule ends" which may attract a new layer of molecules etc., simply by Coulomb forces. Since the acetamino group is n o t a very strong base, an increase in multi
80
layer formation would be expected with decreasing pH. Unfortunately, the pH can n o t be varied very much, for in alkaline solution anodic formation of mercuric oxides occurs, while in strongly acidic solution the substance is hydrolyzed. Nevertheless we found at pH 5.0 surface concentrations even three times greater than at pH 6.6 for solutions 10 -3 M in reactant. The unique adsorption behaviour of the substance studied is also found in a~c. polarography and in galvanostatic pulse measurements which were performed to get differential capacity curves. These curves were in strong contrast to those of related compounds. The phenomena described here are subject to further investigation.
REFERENCES 1 2 3 4 5 6
F.C. Anson and D.A. Payne, J. Electroanal.Chem., 13 (1967) 35. B. Case and F.C. Anson, J. Phys. Chem., 71 (1967) 402. D. Barclay and F.C. Anson, J. Electrochem. Soc., 116 (1969) 438. W. Lorenz and G. Sali~,Z. Phys. Chem. N.F., 29 (1961) 390,408. W. Lorenz and G. Sali~,Z. Phys. Chem., 218 (1962) 259. P. Gr~ndler, in preparation.
77
© Elsevier Sequoia S.A., Lausanne -- Printed in The Netherlands
Preliminary note MULTI LAYER ADSORPTION AT MERCURY ELECTRODES
P. G R U N D L E R and H. WENDLANDT
Karl-Marx-UniversitEt Leipzig, Sektion Chemie, 701 Leipzig, Liebigstr. 18 (D.D.R.) (Received 28th August 1975)
Adsorption processes play an important role in many fields. For practical applications, the surface concentration (surface excess) F of adsorbed species at metal/solution interfaces is of interest. In the course of investigation of some benzenethiosulphonate derivatives a very unusual behaviour of the substance p-acetamino-benzenethiosulphonate (in following shortened to ABTS) was found. NH.CO.CH 3
S02S-
Adsorption of this substance at mercury electrodes was studied by chronocoulometry. Some measurements by a.c.-polarography and galvanostatic pulses were also performed. ABTS is able to form poorly dissociated mercury salts. For substances of this type, Anson and co-workers [1--3] introduced chronocoulometry as a tool for investigation of adsorption processes. Measurements are performed by impressing three potentiostatic voltage steps to the electrode. Primarily, the electrode is held at a potential El (-1.0 V vs. SCE). At this potential, as was indicated by a.c.-polarography, neither adsorption nor any electrochemical reaction occurred. In the first step, the fixed pbtential El is instantaneously switched to a variable potential E2, which permits adsorption, but no electrochemical reaction. Integration of the current flowing during this step gives the charge Q1, which contains double layer charging as well as partial charge transfer (according to Lorenz [4] ). In a second experiment E2 is switched to a fixed potential E3 (+0.2 V vs. SCE). In this step anodic formation of mercury salts occurs. This was proved by d.c.-polarographic studies which established the existence of an anodic polarographic wave with a half wave potential of +0.15 V vs. SCE. Electrolysis currents during this step are integrated and plotted against the square root of
78 time. By extrapolating the straight line of this plot to t = 0 the charge Q2 is obtained. In the same w a y a charge Q3 was determined from a third experiment during which E~ was stepped to E3 directly. From eqn. (1) the a m o u n t of reactant r (mol cm -2) adsorbed on the electrode surface at potential E2 can be calculated, if n (number of electrons involved in anodic salt formation) is known. n F F = Q I (EI-~E2) + Q2 (E2 ~E3 ) - Q3 ( E l s E 3 )
(1)
The value of n was determined b y a coulometric experiment, during which the potential E3 was impressed to a large mercury pool electrode in stirred solution. From the integrated electrolysis current and reactant concentration decrease in the course of electrolysis n was found to be equal 1, i.e. any molecule ABTS was coupled to a one-electron change in anodic salt formation. In our chronocoulometric experiments a simple electronic circuit was used which consisted mainly of an operational amplifier type 741 connected as a potentiostat. An ordinary switch coupled with a small damping capacitor was connected to the non inverting input and allowed to switch in arbitrarily adjustable reference voltages thus producing single pulses at the o u t p u t of the amplifier. The inverting input was connected to the reference electrode (SCE). The o u t p u t of the amplifier was connected through an integrating capacitor to an auxiliary electrode in the cell. Voltage changes of the integrating capacitor were followed b y an oscilloscope, or a recorder in the case of measuring QI. Influence of the offset of the amplifier current was graphically eliminated. Overall response time of the electronic arrangement was 75 ~s. The working electrode was a Kemula type mercury drop electrode with an area of 0.0362 cm 2 . The reference electrode was separated from the cell b y a salt bridge filled with 1 M KNO3. All solutions used were 1 M in KNO3. They also contained a borate (for pH 6.6) or a phosphate (for pH 5.0) buffer system. Values of the adsorption potential E2 varied from - 0 . 8 V to - 0 . 1 V vs. SCE. Concentrations of ABTS lay b e t w e e n 10 -3 M and 10 -2 M. Some chronocoulometric results are shown in Fig. 1. As can be seen from this figure, at potentials more negative than - 6 0 0 mV essentially no adsorpt-ion takes place. From this point to more positive potentials, the adsorbed a m o u n t shows a sharp increase which is followed by a nearly constant region. This is in agreement with a.c. polarograms, which show increasing differential capacities beginning at the same point. In the constant region F depends only on the solution concentration. Further increase in the measured F values in the direction to positive potentials may be due to the beginning of anodic salt formation during m e a s u r e m e n t of QI. The d.c. polarograms indicate slightly rising anodic currents beginning with - 0 . 1 V (for the lowest concentration of ABTS) and shifting to - 0 . 3 V (for the highest reactant concentration). These anodic currents tend to give high results in F. From these measurements as well as from a.c.-polarographic investigations and chronocoulometric studies of other benzenethiosulphonates which do n o t
79
v v
,
/
~' 3C E u E 9 ~ 2O
10
4
f AooJ
/'f f o ~
-10
- O15
° ~ o ~o . ~ r - ~ °
E2//V
vs. SCE
-0
Fig. 1. Surface concentrations of p-acetaminobenzenethiosulphonate solutions. Electrolyte: 1 M K N O s , 0.1 M borate buffer, p H = 6.6. (~) 1 x l0 -3 M, (~) 2 x 10 -3 M, (o) 4 x 10 -3 M, (V) 1 X I 0 -2 M .
contain an acetamino group [6] it is clear, t h a t anions of such compounds are strongly chemisorbed at mercury when potentials are less negative t h a n - 0 . 6 V. In a sharp contrast to the above mentioned studies of other benzenethiosulphonates are the very high values of F f o u n d with ABTS. In the case of benzenethiosulphonate, for instance, F reaches m a x i m u m values of n o t more than 5 x 10 -1° mol cm -2 . This corresponds to a surface area of a b o u t 35 A 2 occupied by each molecule. These data indicate a completely covered electrode surface. As can be seen from Fig. 1, in the case of ABTS, the electrode surface must be completely covered at potentials more positive than - 0 . 6 V already with the lowest reactant concentration studied. This statement holds even if an arrangement of the molecules perpendicular to electrode surface is assumed. For verification we calculated the area covered by one molecule. At reactant concentrations greater than 2 x 10 -3 M, more than one layer of molecules must exist at the electrode surface. We believe t h a t this occurs by a multiple layer adsorption mechanism, for which we propose the following explanation. The thiosulphonate group of the molecule is strongly acidic. Consequently nearly all molecules exist in the form of anions in neutral or slightly acid solution. Anion groups are subject to strong adsorption. On the other hand, any molecule is able to form cations with its acetamino group. One may assume, therefore, that adsorbed molecules form a layer of "cationic molecule ends" which may attract a new layer of molecules etc., simply by Coulomb forces. Since the acetamino group is n o t a very strong base, an increase in multi
80
layer formation would be expected with decreasing pH. Unfortunately, the pH can n o t be varied very much, for in alkaline solution anodic formation of mercuric oxides occurs, while in strongly acidic solution the substance is hydrolyzed. Nevertheless we found at pH 5.0 surface concentrations even three times greater than at pH 6.6 for solutions 10 -3 M in reactant. The unique adsorption behaviour of the substance studied is also found in a~c. polarography and in galvanostatic pulse measurements which were performed to get differential capacity curves. These curves were in strong contrast to those of related compounds. The phenomena described here are subject to further investigation.
REFERENCES 1 2 3 4 5 6
F.C. Anson and D.A. Payne, J. Electroanal.Chem., 13 (1967) 35. B. Case and F.C. Anson, J. Phys. Chem., 71 (1967) 402. D. Barclay and F.C. Anson, J. Electrochem. Soc., 116 (1969) 438. W. Lorenz and G. Sali~,Z. Phys. Chem. N.F., 29 (1961) 390,408. W. Lorenz and G. Sali~,Z. Phys. Chem., 218 (1962) 259. P. Gr~ndler, in preparation.