Study of the Two-Dimensional Phase Formed by Salts of the Cation

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Study of the Two-Dimensional Phase Formed by Salts of the Cation Radical of Heptyl Viologen on Mercury in Aqueous Media J. I. Milla´n, M. Sa´nchez-Maestre, L. Camacho, J. J. Ruiz, and R. Rodrı´guez-Amaro* Department of Physical Chemistry and Applied Thermodynamics, Faculty of Sciences, University of Co´ rdoba, Avda San Alberto Magno s/n, E-14004 Co´ rdoba, Spain Received October 16, 1996. In Final Form: April 21, 1997X In aqueous media, heptyl viologen (HV2+) is reduced via a one-electron transfer to its cation radical (HV•+), the salt of which forms a two-dimensional (2D) phase on a mercury electrode. In this work, we studied the influence of temperature, the reagent concentration, and the type of anion present in the medium on the formation of this 2D phase. The critical temperature for the phase transition was found to be 80-85 °C. The experimental value obtained for the surface area occupied by the heptyl viologen molecule below 30 °C is about 70 Å2; accordingly, the molecules must arrange themselves at an angle with the electrode surface. The anion concentration in a chloride medium was found to exert no appreciable effect on the 2D phase; on the other hand, increasing bromide ion concentrations in the medium shifted the nucleation peak to increasingly more negative potentials.

Introduction 2+),

also known as 1,1′-Disubstituted 4,4′-bipyridyls (V “viologens”, are of great electrochemical interest because they take part in redox reactions that produce stable free radicals which can be used as electrochemical mediators. Viologens can occur in three different oxidation states related by the following equilibria1

V2+ + e- h V•+

(I)

V•+ + e- h V

(II)

where the first reduction step is highly reversible. This process is the subject matter of this work. Specifically, the radical cation formed from heptyl viologen (HV2+) in step I is highly insoluble in water, so reaction I is in fact more complex and involves two steps, namely2,3

HV2+ + e- h HV•+

(IA)

HV•+ + X- h HVXV

(IB)

where step IB involves the binding of the cation radical to anion X- to give an insoluble salt. The precipitate or insoluble film HVX was studied by Jasinski.4 According to this author, the type of anion, pH, and heptyl viologen concentration determine the potential at which the radical is formed and the film adherence, whereas the type of metal substrate (Au, Pt or Ag) influences the kinetics of the process and, also, the film adherence. The rate at which these radical films grow is extremely high and normally controlled by diffusionseven at moderately low potentials. According to Jasinski5 this is a result of the high electron conductivity * Corresponding author. Phone: +34-57-218618. Fax: +34-57218606. E-mail: [email protected]. X Abstract published in Advance ACS Abstracts, June 1, 1997. (1) Bird, C. L.; Kuhn, A. T. Chem. Soc. Rev. 1981, 101, 49. (2) Van Dam, H. T.; Ponjee´, J. J. J. Electrochem. Soc. 1974, 121, 1555. (3) Van Dam, H. T. J. Electrochem. Soc. 1976, 123, 1181. (4) Jasinski, R. J. J. Electrochem. Soc. 1977, 124, 637. (5) Jasinski, R. J. Electrochem. Soc. 1979, 126, 167.

S0743-7463(96)01001-3 CCC: $14.00

of electrodeposited films. The formation of this type of film on optically transparent electrodes has also been studied6,7 and found to exhibit a similar behavior. At potentials more positive than those for process I, a reversible, very sharp peak appears. Scharifker et al.8 designated it X and assigned it to the formation of a twodimensional phase; they found it to be produced on mercury and silver electrodes only. Subsequently, Kitamura et al.9 studied these voltammetric peaks in search for a potential relationship to the alkyl viologen chain. These authors claim that the peaks arise from surface processes. On the basis of studies of Kobayashi et al.10 on bipyridine and methyl viologen, they were ascribed to the following surface electrode reaction:

[HV2+]ads + e- h [HV•+]ads

(III)

As shown in this work, a two-dimensional phase of the salts of the cation radical of heptyl viologen is in fact formed on the mercury electrode. We studied the influence of temperature, the reagent concentration, and the type of anion present in the medium on the phase formation by using cyclic voltammetry and interfacial capacity measurements. Also, we used diagnostic criteria developed in previous work11 for the characterization of this type of process by use of cyclic voltammetry. Experimental Section Practical-grade 1,1′-diheptyl-4,4′-bipyridinium dibromide (purum grade, 97%) was purchased from Aldrich and used without further purification. All other chemicals were Merck, analytical reagent grade, and also used as supplied. Mercury was purified in dilute nitric acid and triply distilled in vacuo. All solutions were made in bidistilled water from a Millipore Milli-Q system and deaerated by bubbling gaseous (6) Bruinik, J.; Kregting, C. G. A. J. Electrochem. Soc. 1978, 125, 1397. (7) Barna, G. G. J. Electrochem. Soc. 1980, 127, 1317. (8) Scharifker, B.; Wehrmann, C. J. Electroanal. Chem. 1985, 185, 93. (9) Kitamura, F.; Ohsaka, T.; Tokuda, K. J. Electroanal. Chem. 1993, 347, 371. (10) Kobayashi, K.; Fuyisaki, F.; Yoshimine, T.; Niki, K. Bull. Chem. Soc. Jpn. 1986, 59, 3715. (11) Sa´nchez Maestre, M.; Rodrı´guez-Amaro, R.; Mun˜oz, E.; Ruiz, J. J.; Camacho, L. J. Electroanal. Chem. 1994, 373, 31-37.

© 1997 American Chemical Society

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Langmuir, Vol. 13, No. 14, 1997 3861

Figure 1. Cyclic voltammogram for 1 mM HV2+ in 0.1 M KBr obtained at v ) 50 mV/s.

Figure 2. Cyclic voltammogram for 1 mM HV2+ in 0.1 M KCl obtained at v ) 50 mV/s.

nitrogen through them. The measuring cell, an Amel 494 model, was thermostated to within (0.1 °C. Unless otherwise noted, the working temperature was 25 °C. Voltammetric measurements were made by using an electronic system consisting of a Model 305 potentiostat and a Model 105 wave generator (both from HQ Instruments), a Norland Prowler digital oscilloscope, and a Houston Instruments 2000 recorder. A hanging mercury drop electrode with a surface area of 2.2 ( 0.05 mm2 was used as the working electrode, and saturated calomel and a platinum wire were employed as the reference and auxiliary electrodes, respectively. A potential where no Faradaic reaction took place was applied for 2 s (delay time) in order to allow the mercury drop to grow and the static mercury drop electrode (SMDE) to stabilize. A potential Ee was then applied over an interval t e = 2 s (equilibrium time), and the cyclic voltammogram was recorded between Ee and a potential E f. Capacity-potential curves were obtained on an Inelecsa Instrument.

if ∆Ep is used to designate the distance between the oxidation and reduction peak potentials and W to denote the peak half-width, then the experimental variations of log ∆Ep and log W with log v are both linear and of slope 0.41 and 0.36, respectively; on the other hand, both ∆Ep and W should be independent of v according to models based on isotherms of the Frumkin or similar types.12 These results are roughly consistent with the expectations for a process involving reduction and the subsequent formation of a two-dimensional (2D) phase of condensed molecules on an electrode as established in the theoretical model developed by our group for 2D phase transitions taking place via nucleation mechanisms in cyclic voltammetry.11 The fundamental equation for the model is

Results and Discussion Figure 1 shows the voltammogram obtained for 1 mM HV2+ in 0.1 M KBr over an Hg electrode at v ) 50 mV/s. The voltammogram exhibits four peaks of which two are cathodic, viz., A1 (Ep = -425 mV) and B (Ep = -600 mV), and two anodic, viz., A 2 (Ep = -415 mV) and K (Ep = -540 mV). In the presence of other anions that are less strongly adsorbed on mercury such as chloride (Figure 2) or sulfate, peaks A1 and A2 appear at more positive potentials (Ep = -380 and -370 mV, respectively), whereas peak B is split into two (B1 at Ep = -650 mV, and B2 at Ep = -660 mV). Peak A1 obtained by cyclic voltammetry is a narrow, sharp peak typical of electrode processes involving molecules immobilized at an electrode. Thus, the peak lacks a diffusion tail; also, the current drops to zero after it. However, its properties are somehow different from the expected results for molecules retained by adsorption on the surface exclusively.12 This peak behaves similarly as those for other viologens.9,13,14 Thus, the variation of log ip with log v for 1 mM HV2+ in a bromide medium at 25 °C is linear, with slope 0.60. On the other hand, when molecules are retained on the electrode by adsorption (Langmuir, Frumkin, and similar isotherms), the plot is also linear but of unity slope.12 Also, (12) Laviron, E. Electroanalytical Chemistry; Bard, A. J., Ed.; Marcel Dekker: New York, 1982; Vol. 12, p 53. (13) Sa´nchez-Maestre, M.; Rodrı´guez-Amaro, R.; Mun˜oz, E.; Ruiz, J. J.; Camacho, L. Langmuir 1994, 10, 723. (14) Salas, R.; Sa´nchez-Maestre, M.; Rodrı´guez-Amaro, R.; Mun˜oz, E.; Ruiz, J. J.; Camacho, L. Langmuir 1995, 11, 1791.

qmb(E - E°)4

i)(

f 5v

[

]

b(E - E°)5

exp (

5f 5v2

(1)

where f ) RT/nF, qm being the overall charge exchanged, E° the standard reduction potential for the phase transition process, and b (V2/s2) a constant related to the nucleation kinetics. The plus signs in the exponential and pre-exponential terms apply to the oxidation process and minus signs to the reduction process. On the basis of eq 1, plots of log ip, log ∆Ep, and log W against log v should all be linear and of slope 0.6, 0.4, and 0.4, respectively. One additional analytical criterion that can be applied to process A involves numerically fitting experimental peaks to eq 1. Figure 3 shows the results of one of the numerical fittings performed for 1 mM HV2+ in the presence of bromide ion at v ) 50 mV/s and T ) 7 °C. The figure compares experimental data obtained after subtraction of the charging current (circles) and the predictions of eq 1 (solid line). The parameter values obtained in the fitting were b ) (8.39 ( 0.12) × 105 mV2/s2 and E° ) -413.5 ( 0.2 mV for the reduction peak and b ) (3.33 ( 0.53) × 106 mV2/s2 and E° ) -417.8 ( 0.5 mV for the oxidation peak. As can be seen, the theoretical predictions were quite consistent with the experimental data. Equation 1 predicts that, at E ) E°, the current must be zero, both for the reduction and for the oxidation process (i.e., a 2D phase transition process must go through a node or zone of zero Faradaic current between the reduction and oxidation peaks). However, as can be seen in Figure 3, while the E° values obtained in the fittings

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Figure 3. Cyclic voltammogram for 1 mM HV2+ in 0.1 M KBr obtained at T ) 7 °C and v ) 50 mV/s, after subtraction of the charging current (circles). The solid line corresponds to the predictions of eq 1 for b ) 8.93 × 105 mV2/s2 and E°) -413.5 mV (reduction peak), and b ) 3.33 × 106 mV2/s2 and E°) -417.8 mV (oxidation peak).

Figure 4. Variation of the charge of peak A1 with temperature for 1 mM HV2+ in 0.1 M KBr at v ) 50 mV/s.

of the reduction and oxidation processes differ by only about 4 mV, the current does not fall to zero between the two peaks. This was also the case at other temperatures where fittings were made. One plausible explanation for this phenomenon, previously exposed elsewhere,15 involves the presence of the anion in the 2D phase; if such were the case, the anion might alter the E° value for the 2D phase formation relative to its destruction. Therefore, as suggested by Scharifker et al.8 and shown above, peaks A1 and A2 correspond to a 2D phase transition process. Integration with respect to time of the peak for process A (whether the reduction or oxidation peak) allowed us to calculate the charge Q exchanged during the formation or disappearance of the 2D phase. Figure 4 shows the Q values obtained at a variable temperature. Below 30 °C, Q was ca. 23 µC/cm2, so the area occupied by HV‚+ should (15) Sa´nchez-Maestre, M.; Rodrı´guez-Amaro, R.; Mun˜oz, E.; Ruiz, J. J.; Camacho, L. J. Electroanal. Chem. 1995, 390, 21-27.

Milla´ n et al.

Figure 5. Plots of Ep vs log [X-] for peaks A1 (b) and B (O).

be about 70 Å2/molecule. Because the HV2+ occupies a surface area of about 100 Å 2/molecule parallel to the electrode and about 20 Å2 normal to it,8 the experimental value, 70 Å2, suggests that the molecule is tilted relative to the electrode surface, similarly as 4,4′-bipyridine13 and methyl viologen.14 Unfortunately, no specific model for the situation of the heptyl viologen molecule relative to the electrode can be established owing to the different possible conformations that the alkyl chains of the molecule can take.16 Above 50 °C, Q drops sharply (Figure 4). Extrapolation of Q to zero leads to a temperature of about 82 °C; however, peaks A1 and A2 were never observed above 77 °C. The disappearance of these peaks was a result of the 2D phase formed by the cation radical bromide being destroyed;14,17 therefore, the corresponding critical phase transition temperature, Tc, must lie in the range 80-85 °C. The Q values for peaks A1 and A2 in the chloride medium were similar to those obtained in the bromide medium. We also examined the influence of the anion in the supporting electrolyte on peak A1 (Figure 5). In the chloride medium, Ep was virtually independent of the chloride concentration, at least over the concentration range studied; by contrast, in the bromide medium, Ep was found to depend strongly on the anion concentration. Thus, the peak was shifted to more negative values (by about -26 mV/dec) with increase in the concentration of bromide ion. This seems to contradict the potential involvement of the anion in the 2D phase since an increase in the anion concentration should favor the formation. However, it should be noted that the process is also governed by a term that takes account of the energy required to displace the anions adsorbed on the electrode. Thus, with chloridesa weakly adsorbed anionsthe two effects appear to offset each other; on the other hand, the strong adsorption of bromide ion on the electrode surface favors the second effect, so the more concentrated the anion is, the more difficult it is for the 2D phase to be formed. Regarding peak B, the variation of the peak potential for the process with the logarithm of the bromide (16) Cotton, T. M.; Kim, J. H.; Uphaus, R. A. Microchem. J. 1990, 42, 44. (17) Stenina, E. V.; Damaskin, B. B. J. Electroanal. Chem. 1993, 349, 31.

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Langmuir, Vol. 13, No. 14, 1997 3863

Figure 6. Plots of Ep vs T (°C) for peaks A1 (b) and B (O). All conditions as in Figure 4.

concentration was linear (Figure 5) and of slope 44 mV/ dec, i.e., the process was favored by increasing concentrations of the anion in the medium. This, together with the presence of peak K in the anodic scan, which was identified elsewhere as belonging to the stripping of a precipitate,6,18 confirms that peak B is the result of two consecutive processes, viz., the one-electron reduction of the molecules that reach the electrode by diffusion (process IA) and the precipitation of the cation radical formed in this process with the anion present in the medium as counterion (process IB). In the chloride medium (Figure 5), peaks B1 and B 2 also vary linearly with the anion concentration, with a slope of 28.5 and 67 mV/dec, respectively. The appearance of these peaks at potentials more negative than those obtained in the bromide medium appears to be related to a higher solubility of the precipitate formed in the chloride medium, the formation of which is also favored by increasing concentrations of the halide. Figure 6 shows the variation of the potentials corresponding to peaks A1 and B with temperature. As can be seen, both varied similarly. On the other hand, the peak potentials for processes A and B in the bromide medium varied linearly with the heptyl viologen concentration, with a slope of 18 and 54 mV/dec, respectively. Figure 7.1 shows the differential capacity-potential curves obtained for a 0.1 M KBr solution at pH 6 in the absence of HV2+ (dashed line) and the presence of 1 mM HV2+ (solid line), at 300 Hz, ∆E ) 10 mV, and v ) 2 mV/s. Figure 7.2 shows a voltammogram (forward scan only) obtained at the same concentration and v ) 50 mV/s. At potentials more positive than those of appearance of peak A1 (zone I), the capacity is somewhat higher than that for the supporting electrolyte; this phenomenon was previously observed in other viologens.10 However, the appearance of peak A1 causes an abrupt drop in the capacity that can be related to the formation of the 2D phase of the radical cation bromide HV•+Br-, similarly as for 4,4′-bipyridine.13 Although the experimental capacity curves were obtained in parallel with Faradaic processes, the voltammetric peak A1 (Figure 7.2) does not give rise to a significant peak in the C-E curves. As in 4,4′bipyridin,13 this must be related to the prevalence of the capacitive phenomenonsat least at the temperature where the recording shown in Figure 7 was obtained. Beyond peak A1 (zone II), where the capacity is minimum, this parameter rises to values exceeding those for the supporting electrolyte (zone III). The rise cannot be attributed (18) Bewick, A.; Lowe, A. C.; Wederell, C. W. Electrochim. Acta 1983, 28, 1899.

Figure 7. Plot of C (1) and the current (2) as a function of E. C was determined at a frequency of 300 Hz and a scan rate of 2 mV/s by using a 10 mV pulse. The dashed line corresponds to a solution of 0.1 M of KBr containing no HV2+; the solid line corresponds to the same solution containing 1 mM HV2+. The voltammogram was obtained from the same HV2+ solution at v ) 50 mV/s.

Figure 8. Plots of C vs E at a variable concentration of HV2+: (a) 0.05 mM, (b) 0.1 mM, (c) 0.5 mM, and (d) 1 mM. All others conditions as in Figure 7.

to a Faradaic phenomenon (see Figure 2); rather, it might be related to a loss of order in the 2D phase. In this zone, the capacity peaks at potentials close to -550 mV. Then, it falls at the potentials corresponding to the appearance of peak B (zone IV). It should be noted that, as with peak A1, no matching peak was observed in the C-E curves despite the fact that a Faradaic process took place in parallel. Figure 8 shows the variation of the interfacial capacity as a function of the concentration of heptyl viologen in solution. As can be seen, the drop corresponding to the appearance of peak A1 shifts to more positive potentials as the concentration is raised, similarly as peak A1 in voltammetry. In addition, the width of zone II in Figure 7.2 increases with increasing concentrationsconcomitantly with an increase in the capacity. On the other hand, the potential of appearance of the maximum (peak) obtained

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Milla´ n et al.

Figure 9. Plots of C vs E at a variable temperature: (a) 2 °C, (b) 11 °C, (c) 25 °C, and (d) 51 °C. All others conditions as in Figure 7.

in zone III varies very little with the concentration. By contrast, the subsequent drop, zone IV, is indeed affected by this variable: it is shifted to less negative potentials as the concentration is raised, similarly as peak B in voltammetry. As can be seen from Figure 9, which shows the variation of the interfacial capacity with temperature, a peak is obtained at the intersect of zones I and II that increases with increase in the temperature. Such a peak must be

related to the Faradaic contribution to the overall signal, the magnitude of which must increase with increasing temperature (i.e., with increasing disorder in the monolayer). In addition, this capacity peak is shifted to more negative values as T is increased, consistent with the shift in the potential of the voltammetric peak A1 (Figure 6). Finally, the second drop (zone IV) disappears at about 74 °C; this suggests that, above such a temperature, the product formed at these potentials is more soluble so it

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Langmuir, Vol. 13, No. 14, 1997 3865

and IV was higher in the former case as a result of the shift to more positive potentials in peak A1 and to more negative ones in peaks B1 and B2. It should be noted that, as in the bromide medium, the capacity maximum observed in zone III was also obtained at about -550 mV. Although the overall process that gives peak B, described above, is split into peaks B1 and B2, these must also correspond to the formation of an insoluble precipitate unless the first monolayer (peak B1) is formed at potentials slightly more positive than those for the three-dimensional precipitation.

Figure 10. Plot of C (1) and the current (2) as a function of E. The dashed line corresponds to a solution of 0.1 M of KCl containing no HV2+, while the solid line corresponds to a similar solution containing 1 mM HV2+. All others conditions as in Figure 7.

does not precipitatesand hence causes no abrupt change in the interfacial capacity. The results obtained with chloride as the counterion (Figure 10) were quite similar to those for bromide with the exception that the potential range between zones I

Conclusions As shown above, the salts of the cation radical of heptyl viologen form a two-dimensional phase on a mercury electrode in an aqueous medium. In such a phase, molecules lie at an angle to the electrode surface (the experimental surface area is 70 Å2 below 30 °C). The formation of this 2D phase is inhibited by the presence of anions that are strongly adsorbed on the electrode (e.g., bromide). On the other hand, less strongly adsorbed anions such as chloride have little influence on the phenomenon. After the 2D film is formed, the molecules rearrange (possibly with some loss of order, as suggested by the capacity-potential curves). However, as reflected in the cyclic voltammograms, the phenomenon is quite reversible. Acknowledgment. The authors wish to express their gratitude to Spain’s DGICyT for financial support of this research in the framework of Projects PB94-0446 and PB94-0448. LA9610011