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Mechanism of dihydroxyfumaric acid oxidation on a mercury electrode
Electraehimia,
Act,,.
Vol. 25, pp. 833-837.
Pergamon Press Ltd. 1980. Printed in Great Britrun.
MECHANISM OF DIHYDROXYFUMARIC ACID OXIDATION ON A MERCURY ELECTRODE M. DOM~NGUEZ and E. VALERA Departamento de Quhnica Fisica, Facultad de Ciencias, Universidad de Sevilla, Spain (Received 29 January 1979) Abstmct A study was carried out on the oxidation mechanism ofdihydroxyfumaric acid over DME. The two-electron transfer takes place through a post-kinetic mechanism. The overall reaction was determined together with the kinetic parameters at the foot ofthe wave. A basic plan was drawn up showing experimental and theoretical data.
INTRODUCTION
Although several studies have been carried out on Lascorbic acid oxidation[ l-31 and u-araboascorbic acid oxidation[4,5] (two ene-diol compounds), the mechanism over DME of dihydroxyfumaric acid, also an ene-diol, has not been discussed, as far as the authors can tell, in any published work. The present study is based on an analysis of the i-E curves for dihydroxyfumaric acid at varying pH vahtes. An analysis was also carried out on the influence of the temperature and velocity of the potential sweep on the i-E curves and the concentration of the electroactive species. The products from the oxidation process and the total number of electrons that intervene are also determined. With these results, together with the calculation of the Tafel slopes and reaction orders, a scheme is proposed for the overall reaction of the oxidation process together with the mechanism for potentials corresponding to the foot of the wave.
thermostatted and possessed the corresponding inlets for the three electrodes, a nitrogen flow and thermometer. The reference electrode was one of saturated calomel. The working electrode used for the polarographic measurements was a mercury drop constructed with a B400 Radiometer capillary with the followingcharacteristics:m= 13OOmgs-‘andt=6.35sin our but&red solution at pH = 7, h = 4Ocm and open circuit. In the processes for obtaining the reaction products, the working electrode was a mercury pool. The auxiliary electrode was the typical platinum one which was also used in the voltametric measurements. In
EXPERIMENTAL
Apparatus The polarographic curves were registered automatically by means of a PO4 Radiometer where the damping circuit had been completely suppressed. A Beckman research potentiometer was used to measure the start potentials as well as the pH. The i-E curves were traced point to point with the aid of the galvanometer and the power source of the polarograph ; the potential being measured, in all instances, by the Beckman potentiometer. An Amel triangular wave generator was used for the voltametric techniques, together with an Amel potentiostat with a built in correction for ohmic drop, an X-Y recorder from Hewlett-Packard and an oscilloscope from Amel. The spectrometer used was a Beckman DB-GT model for ultraviolet. Cells and electrodes The polarographic measurements were made in Radiometer V519 cell with a thermostat. The cell used for the microcoulometric measurements was the classic 2 ml one, whilst that used for the voltametric measurements with a capacity of approx. 50m1, was
Fig. 1. Variation of i, with time. (a) pH = 2.90. (b) pH = 3.90.
(c) pH = 4.90.(d) pH = 5.90.(e)pH = 6.50.(f) pH = 6.90.(g) pH = 7.80. (h) pH = 9.70. (i) pH = 10.90.(j) pH = 11.90. I333
M. DOMfNGUEZ
834
these last measurements the working electrode was a hanging mercury electrode (Radiometer B404 model). Solutions, products and measurements All products used were Merck p-a. except the disodium diketosuccinate which was obtained by the method of Tyurenkova and Lipunova[6]. As a supporting electrolyte solution, a buffered solution was used with the following components and-concentrations : 0.04 M acetic acid, 0.04 M phosphoric acid, 0.04 M ammonium nitrate, and 0.2 M NaOH, which were mixed in varying proportions according to the pH desired. The ionic strength was adjusted with NaNO, to 0.2 M. Owing to the instability of dihydroxyfumaric acid solutions these were prepared immediately prior to each experiment and were added to the supporting solution with the oxygen removed. All measurements were taken in a nitrogen medium at 25 + O.l”C, with the exception of those cases where the effect of temperature char.ge was studied. To avoid depleting the solution, those ‘measurements for plotting the polarographic curve point by point were taken with each first drop, manually synchronfzing the application of the potential with the fall of the drop. The Tafel curves were taken on the rising portion of the polarographic wave. All values of i or E were corrected for the residual current and the iR drop. RESULTS
General behauiour The dihydroxyfumaric
I
acid oxidises on the DME,
I
9
”
I
Fig. 2. Variation of E,,, with PH. c = 5 x IO-* M,
AND E. VALERA
producing one polarographic curve in the pH range ca j-10. The oxidation wave varies with time due to the instability of the dihydroxyfumaric acid with respect to its self-oxidation even in an inert atmosphere. For this reason, we have made a polarographic study of the instability of diiydroxyfumaric acid. Figure 1 gives the plot of i, us t at varying pH values. It is evident that the dihydroxyfumaric acid is very unstable in acid and basic mediums, being relatively stable in a neutral medium. Therefore, we restrict the study of the oxidation mechanism to neutral medium. Microcoulometric measurements The number of electrons taking part in the reaction was typically calculated from the decrease of the limiting current with the time of electrolysis. The measurements were corrected for the instability of dihydroxyfumaric acid. The average value obtained was 2.0. Polarographic behaviour The limiting current is independent of the pH, whilst El/2 shifts to more negative potentials with an increase in pH (Fig. 2). The values of the slopes of the linear segments are approx. -59mVand -28mVperpH unit. The logarithmic analysis of the polarographic wave gave straight lines with a slope of 3OmV per decade that are practically independent of the pH. The limiting current is controlled by the diffusion : log i, us log h,,, is linear with an approximate slope of 0.5 and a temperature coefficient of 1.0 % per “C, and is proportional to the concentration of the dihydroxyfumaric acid in the bulk of the solution.
Fig. 3. Variation of i with h”“. pH = 7.00. Applied potential: (a)0.053 V. (b) - 0.02OV. (c) - 0.098 V. (d) - 0.128 V. (f) - 0.149 v.
Mechanism of dihydroxyfumaric acid oxidation
Fig. 4. Representation of Tafet’s law. (a] pH = 7.40. (b) pH = 7.32, (c) pH = 7.14. (d) pH = 6.92. (e) pH = 6.82. (f) pH = 6.56. (g) pH = 6.40.
Fig.
Et,, is independent of the dihydroxytumaric acid concentration, but varies in a linear way with log td with an approximate slope of - 16 mV. The temperature coefficient of Et,, is 1 mV per “C. Figure 3 illustrates the influence of the height of the mercury column at maximum current at varying potentials in the wave. The potentials - 149, - 140 and - 128 mV correspond to that zone where the condition (i,&) +z 0.02 is satisfied. A potential of 53 mV corresponds to the zone (i,& > 0.8.
Electrolysis and identiJication of compounds
Tafel curves and reaction orders
(a) Iajluence ofpH. The Tafel slope is independent of the pH with an average value of 39 mV per decade (Fig. 4). The reaction order with respect to the H+ concentration is - 1.1 in the pH range 6.4-7.4. The order is independent of the potential in the area where Tafel’s law is obeyed, as shown in Fig. 5. (b) Influence of dihydroxyfumaric acid concentration. The reaction order with respect to the concentration of dihydroxyfumaric acid is 1 and independent of the pH and of the potential at which it is measured. Influence of potential sweep velocity At normal concentrations of dihydroxyfumaric acid, no reduction voltametric wave is observed, even with the maximum sweep velocity reached with our system (4 V/s). The oxidation wave displaces towards more anodic potential values when the sweep velocity is increased, Fi ures 6 and 7 give the plots of E, us logu and i, us vi ?2
5.
Reaction
835
orders.
(a)
E = -O.lSOV.
E = - 0.19OV. (c) E = - 0.2OOV.(d) E = - 0.2tOV.
(b) (e)
E = - 0.22ov.
The oxidation product of dihydroxyfumaric acid was identified by ultraviolet spectroscopy of a solution OFdihydroxyfumaric acid at pH 7 which had undergone electrolysis at a controlled potential. The oxidation product was identified as disodium diketosuccinate because of its ,I,.. characteristics (230 nm). In Fig. 8 the spectrum of the electrolysis product after eight hours is shown. The characteristic band of dihydroxyfumaric acid (approx. 280nm) and the weak band of diketosuccinate can be observed. This last product is very unstable in aqueous solutions, and therefore we have not studied its electrochemical reduction. DISCUSSION Overall
reaction
The data collected demonstrate that the oxidation wave of dihydroxyfumaric acid is produced by a polarographically reversible two-electron transfer, in which two protons intervene in the interval pH 6.4-7.4 (slope J&z us PH - 59mV decade-t). However, the peculiar variation of i,,, OS h”’ confirms that the process under study is of the post-kinetic type, with a rapid chemical reaction. This Fact is confirmed by the variation of E,,, with logt,[l] ; the independence of Et,, with the concentration supposes a first order or pseudo first order reaction with respect to the electrodic product For the kinetic process[7]. Thus, the overall reaction in the neutral medium would be that of Fig. 9.
M.
836
DOM~NGUEZANLIE. VALERA
Fig. 6. Variation of E, with potential sweep velocity. pH = 7.00 ‘OOC,
,c=c\ HO
IOH
.
Hz0
*
-
2r
coo-
. 2t.j’
Fig. 9. Overall reaction. Our voltametric data confirm this hypothesis. For this type of process and for low sweep velocities, the peak potential variation is in agreement with[g] :
E, = E,,, -
~(0.780 - In Jk,la)
where k, is the velocity constant of the hydration reaction and a = nFv/RT The fact that the peak potential shifts towards more anodic values with an increase in the sweep velocity is in agreement with the same equation. Oxidation mechanism at the footof the wase With the at&rnentioned conclusions, the mechanism can be represented by tbe following reaction kinetic pathways, as shown in Fig. 10 where reaction 1 represents the dissociation of H+ from carbon 2, and reactions 2 and 3 represent the one-electron transfers with presumably formation of a free radical. Reaction 4 represents the hydration of the carbonyl group. Assuming reaction 3 as the rds and applying the approximation of the equilibrium state, the value of the anodiccurrent intensity on the rising portion of the
,..I’p of i,, with potential sweep velocity. pH = 7.00.
Fig. 7. Variation
I 110
Fig. 8. Ultraviolet
1sLl
110
spectrum of dihydroxyfumaric
150
190
acid electrolysis product.
xc
Mechanism of dihydroxyfumaric acid oxidation
1.
-ooc, /OH c=c,
L
‘OOC,
COO-T
HO’
c,C/O-.
wave can be expressed (taking potential to be negligible) as :
H’
‘coo-
Hd
837
i = 2FK,K2k3 &$-
n
._
-ooc, 4.
0
c-c<
O@
*H*O
coo-
k4
O\H /OH
-OOc\
c-c,
=e
k-4
o”
coo-
the electrokinetic
expC(2 -
1
#WJY~~~l
where CT is the total concentration of dibydroxyfumaric acid in the bulk of the solution; CH is the concentration of the H+ ions; KI and Kz are the equilibrium constants of reactions 1 and 2 and k-, is the rate constant of the rate controlling step. The comparison between the theoretical values deduced in the aforementioned sequences and the experimental results is shown in Table 1 (where fl is assumed to have an approximate value of 0.5). REFERENCES
Fig. 10. Reaction kinetic pathway of dihydroxyfumaric acid
2.
Table 1.
Parameter Tafel slope Order with respect to dihydroxyfumaric acid concentration C$d;nwith respect to
1.
Value
3.
Theoretical (B E 0.5) Experimental
4. 5.
39 mV
39mV 6.
1 -1
1
I
- 1.1
8
D. H. M. Kern, J. Am.ch.?m.Sot. 76,1011(1954); 75,2473 (1953). S. Ono. M. Takaei and T. Wasa. &I[. them. Sot. Jaoan 31,358 (1958). I. J. Ruiz, A. Aldax and M. Dominguez, Can. J. Chem. 55, 2799 (1977); 56, 1533 (1978). M. Dominguex, A. Aldax and F. Sanchez-Burgos, J. electroanal. Chem. 68, 345 (1976). M. Dominguez, A. Aldaz and F. Shnchex-Burgos, An. Quim. Farm. 74, 199 (1978). G. N. Tyurenkova and G. N. Lipunova, Zh. Prikl. K&m., Len+. 41, 664 (1968). R. Guide& Elecrrounnlyrical Chemistry Vol. 5, pp. 149-374. Marcel Dekker, New York (1971). R. S. Nicholson and 1. Shain, Anolyt. Chem. 37, 191 (1965); 37, 179 (1965).
Act,,.
Vol. 25, pp. 833-837.
Pergamon Press Ltd. 1980. Printed in Great Britrun.
MECHANISM OF DIHYDROXYFUMARIC ACID OXIDATION ON A MERCURY ELECTRODE M. DOM~NGUEZ and E. VALERA Departamento de Quhnica Fisica, Facultad de Ciencias, Universidad de Sevilla, Spain (Received 29 January 1979) Abstmct A study was carried out on the oxidation mechanism ofdihydroxyfumaric acid over DME. The two-electron transfer takes place through a post-kinetic mechanism. The overall reaction was determined together with the kinetic parameters at the foot ofthe wave. A basic plan was drawn up showing experimental and theoretical data.
INTRODUCTION
Although several studies have been carried out on Lascorbic acid oxidation[ l-31 and u-araboascorbic acid oxidation[4,5] (two ene-diol compounds), the mechanism over DME of dihydroxyfumaric acid, also an ene-diol, has not been discussed, as far as the authors can tell, in any published work. The present study is based on an analysis of the i-E curves for dihydroxyfumaric acid at varying pH vahtes. An analysis was also carried out on the influence of the temperature and velocity of the potential sweep on the i-E curves and the concentration of the electroactive species. The products from the oxidation process and the total number of electrons that intervene are also determined. With these results, together with the calculation of the Tafel slopes and reaction orders, a scheme is proposed for the overall reaction of the oxidation process together with the mechanism for potentials corresponding to the foot of the wave.
thermostatted and possessed the corresponding inlets for the three electrodes, a nitrogen flow and thermometer. The reference electrode was one of saturated calomel. The working electrode used for the polarographic measurements was a mercury drop constructed with a B400 Radiometer capillary with the followingcharacteristics:m= 13OOmgs-‘andt=6.35sin our but&red solution at pH = 7, h = 4Ocm and open circuit. In the processes for obtaining the reaction products, the working electrode was a mercury pool. The auxiliary electrode was the typical platinum one which was also used in the voltametric measurements. In
EXPERIMENTAL
Apparatus The polarographic curves were registered automatically by means of a PO4 Radiometer where the damping circuit had been completely suppressed. A Beckman research potentiometer was used to measure the start potentials as well as the pH. The i-E curves were traced point to point with the aid of the galvanometer and the power source of the polarograph ; the potential being measured, in all instances, by the Beckman potentiometer. An Amel triangular wave generator was used for the voltametric techniques, together with an Amel potentiostat with a built in correction for ohmic drop, an X-Y recorder from Hewlett-Packard and an oscilloscope from Amel. The spectrometer used was a Beckman DB-GT model for ultraviolet. Cells and electrodes The polarographic measurements were made in Radiometer V519 cell with a thermostat. The cell used for the microcoulometric measurements was the classic 2 ml one, whilst that used for the voltametric measurements with a capacity of approx. 50m1, was
Fig. 1. Variation of i, with time. (a) pH = 2.90. (b) pH = 3.90.
(c) pH = 4.90.(d) pH = 5.90.(e)pH = 6.50.(f) pH = 6.90.(g) pH = 7.80. (h) pH = 9.70. (i) pH = 10.90.(j) pH = 11.90. I333
M. DOMfNGUEZ
834
these last measurements the working electrode was a hanging mercury electrode (Radiometer B404 model). Solutions, products and measurements All products used were Merck p-a. except the disodium diketosuccinate which was obtained by the method of Tyurenkova and Lipunova[6]. As a supporting electrolyte solution, a buffered solution was used with the following components and-concentrations : 0.04 M acetic acid, 0.04 M phosphoric acid, 0.04 M ammonium nitrate, and 0.2 M NaOH, which were mixed in varying proportions according to the pH desired. The ionic strength was adjusted with NaNO, to 0.2 M. Owing to the instability of dihydroxyfumaric acid solutions these were prepared immediately prior to each experiment and were added to the supporting solution with the oxygen removed. All measurements were taken in a nitrogen medium at 25 + O.l”C, with the exception of those cases where the effect of temperature char.ge was studied. To avoid depleting the solution, those ‘measurements for plotting the polarographic curve point by point were taken with each first drop, manually synchronfzing the application of the potential with the fall of the drop. The Tafel curves were taken on the rising portion of the polarographic wave. All values of i or E were corrected for the residual current and the iR drop. RESULTS
General behauiour The dihydroxyfumaric
I
acid oxidises on the DME,
I
9
”
I
Fig. 2. Variation of E,,, with PH. c = 5 x IO-* M,
AND E. VALERA
producing one polarographic curve in the pH range ca j-10. The oxidation wave varies with time due to the instability of the dihydroxyfumaric acid with respect to its self-oxidation even in an inert atmosphere. For this reason, we have made a polarographic study of the instability of diiydroxyfumaric acid. Figure 1 gives the plot of i, us t at varying pH values. It is evident that the dihydroxyfumaric acid is very unstable in acid and basic mediums, being relatively stable in a neutral medium. Therefore, we restrict the study of the oxidation mechanism to neutral medium. Microcoulometric measurements The number of electrons taking part in the reaction was typically calculated from the decrease of the limiting current with the time of electrolysis. The measurements were corrected for the instability of dihydroxyfumaric acid. The average value obtained was 2.0. Polarographic behaviour The limiting current is independent of the pH, whilst El/2 shifts to more negative potentials with an increase in pH (Fig. 2). The values of the slopes of the linear segments are approx. -59mVand -28mVperpH unit. The logarithmic analysis of the polarographic wave gave straight lines with a slope of 3OmV per decade that are practically independent of the pH. The limiting current is controlled by the diffusion : log i, us log h,,, is linear with an approximate slope of 0.5 and a temperature coefficient of 1.0 % per “C, and is proportional to the concentration of the dihydroxyfumaric acid in the bulk of the solution.
Fig. 3. Variation of i with h”“. pH = 7.00. Applied potential: (a)0.053 V. (b) - 0.02OV. (c) - 0.098 V. (d) - 0.128 V. (f) - 0.149 v.
Mechanism of dihydroxyfumaric acid oxidation
Fig. 4. Representation of Tafet’s law. (a] pH = 7.40. (b) pH = 7.32, (c) pH = 7.14. (d) pH = 6.92. (e) pH = 6.82. (f) pH = 6.56. (g) pH = 6.40.
Fig.
Et,, is independent of the dihydroxytumaric acid concentration, but varies in a linear way with log td with an approximate slope of - 16 mV. The temperature coefficient of Et,, is 1 mV per “C. Figure 3 illustrates the influence of the height of the mercury column at maximum current at varying potentials in the wave. The potentials - 149, - 140 and - 128 mV correspond to that zone where the condition (i,&) +z 0.02 is satisfied. A potential of 53 mV corresponds to the zone (i,& > 0.8.
Electrolysis and identiJication of compounds
Tafel curves and reaction orders
(a) Iajluence ofpH. The Tafel slope is independent of the pH with an average value of 39 mV per decade (Fig. 4). The reaction order with respect to the H+ concentration is - 1.1 in the pH range 6.4-7.4. The order is independent of the potential in the area where Tafel’s law is obeyed, as shown in Fig. 5. (b) Influence of dihydroxyfumaric acid concentration. The reaction order with respect to the concentration of dihydroxyfumaric acid is 1 and independent of the pH and of the potential at which it is measured. Influence of potential sweep velocity At normal concentrations of dihydroxyfumaric acid, no reduction voltametric wave is observed, even with the maximum sweep velocity reached with our system (4 V/s). The oxidation wave displaces towards more anodic potential values when the sweep velocity is increased, Fi ures 6 and 7 give the plots of E, us logu and i, us vi ?2
5.
Reaction
835
orders.
(a)
E = -O.lSOV.
E = - 0.19OV. (c) E = - 0.2OOV.(d) E = - 0.2tOV.
(b) (e)
E = - 0.22ov.
The oxidation product of dihydroxyfumaric acid was identified by ultraviolet spectroscopy of a solution OFdihydroxyfumaric acid at pH 7 which had undergone electrolysis at a controlled potential. The oxidation product was identified as disodium diketosuccinate because of its ,I,.. characteristics (230 nm). In Fig. 8 the spectrum of the electrolysis product after eight hours is shown. The characteristic band of dihydroxyfumaric acid (approx. 280nm) and the weak band of diketosuccinate can be observed. This last product is very unstable in aqueous solutions, and therefore we have not studied its electrochemical reduction. DISCUSSION Overall
reaction
The data collected demonstrate that the oxidation wave of dihydroxyfumaric acid is produced by a polarographically reversible two-electron transfer, in which two protons intervene in the interval pH 6.4-7.4 (slope J&z us PH - 59mV decade-t). However, the peculiar variation of i,,, OS h”’ confirms that the process under study is of the post-kinetic type, with a rapid chemical reaction. This Fact is confirmed by the variation of E,,, with logt,[l] ; the independence of Et,, with the concentration supposes a first order or pseudo first order reaction with respect to the electrodic product For the kinetic process[7]. Thus, the overall reaction in the neutral medium would be that of Fig. 9.
M.
836
DOM~NGUEZANLIE. VALERA
Fig. 6. Variation of E, with potential sweep velocity. pH = 7.00 ‘OOC,
,c=c\ HO
IOH
.
Hz0
*
-
2r
coo-
. 2t.j’
Fig. 9. Overall reaction. Our voltametric data confirm this hypothesis. For this type of process and for low sweep velocities, the peak potential variation is in agreement with[g] :
E, = E,,, -
~(0.780 - In Jk,la)
where k, is the velocity constant of the hydration reaction and a = nFv/RT The fact that the peak potential shifts towards more anodic values with an increase in the sweep velocity is in agreement with the same equation. Oxidation mechanism at the footof the wase With the at&rnentioned conclusions, the mechanism can be represented by tbe following reaction kinetic pathways, as shown in Fig. 10 where reaction 1 represents the dissociation of H+ from carbon 2, and reactions 2 and 3 represent the one-electron transfers with presumably formation of a free radical. Reaction 4 represents the hydration of the carbonyl group. Assuming reaction 3 as the rds and applying the approximation of the equilibrium state, the value of the anodiccurrent intensity on the rising portion of the
,..I’p of i,, with potential sweep velocity. pH = 7.00.
Fig. 7. Variation
I 110
Fig. 8. Ultraviolet
1sLl
110
spectrum of dihydroxyfumaric
150
190
acid electrolysis product.
xc
Mechanism of dihydroxyfumaric acid oxidation
1.
-ooc, /OH c=c,
L
‘OOC,
COO-T
HO’
c,C/O-.
wave can be expressed (taking potential to be negligible) as :
H’
‘coo-
Hd
837
i = 2FK,K2k3 &$-
n
._
-ooc, 4.
0
c-c<
O@
*H*O
coo-
k4
O\H /OH
-OOc\
c-c,
=e
k-4
o”
coo-
the electrokinetic
expC(2 -
1
#WJY~~~l
where CT is the total concentration of dibydroxyfumaric acid in the bulk of the solution; CH is the concentration of the H+ ions; KI and Kz are the equilibrium constants of reactions 1 and 2 and k-, is the rate constant of the rate controlling step. The comparison between the theoretical values deduced in the aforementioned sequences and the experimental results is shown in Table 1 (where fl is assumed to have an approximate value of 0.5). REFERENCES
Fig. 10. Reaction kinetic pathway of dihydroxyfumaric acid
2.
Table 1.
Parameter Tafel slope Order with respect to dihydroxyfumaric acid concentration C$d;nwith respect to
1.
Value
3.
Theoretical (B E 0.5) Experimental
4. 5.
39 mV
39mV 6.
1 -1
1
I
- 1.1
8
D. H. M. Kern, J. Am.ch.?m.Sot. 76,1011(1954); 75,2473 (1953). S. Ono. M. Takaei and T. Wasa. &I[. them. Sot. Jaoan 31,358 (1958). I. J. Ruiz, A. Aldax and M. Dominguez, Can. J. Chem. 55, 2799 (1977); 56, 1533 (1978). M. Dominguex, A. Aldax and F. Sanchez-Burgos, J. electroanal. Chem. 68, 345 (1976). M. Dominguez, A. Aldaz and F. Shnchex-Burgos, An. Quim. Farm. 74, 199 (1978). G. N. Tyurenkova and G. N. Lipunova, Zh. Prikl. K&m., Len+. 41, 664 (1968). R. Guide& Elecrrounnlyrical Chemistry Vol. 5, pp. 149-374. Marcel Dekker, New York (1971). R. S. Nicholson and 1. Shain, Anolyt. Chem. 37, 191 (1965); 37, 179 (1965).