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Electrochemical behavior of some triazine derivatives at glassy carbon electrode in non-aqueous media
Journal of
Electroanalytical Chemistry Journal of Electroanalytical Chemistry 580 (2005) 245–254 www.elsevier.com/locate/jelechem
Electrochemical behavior of some triazine derivatives at glassy carbon electrode in non-aqueous media N. Farzinnejad
a,*
, A.A. Miran Beigi
b,*
, L. Fotouhi a, K. Torkestani b, H.A. Ghadirian
b
a
b
Department of Chemistry, Faculty of Science, Al-Zahra University, P.O. Box 1993891176, Tehran, Iran Standard Laboratory, Physical and Chemical Analyses Group, Research Institute of Petroleum Industry, P.O. Box 18745-4163, Tehran, Iran Received 12 October 2004; received in revised form 4 March 2005; accepted 17 March 2005 Available online 23 May 2005
Abstract Electroreduction of 4-amino-6-methyl-3-thio-1,2,4-triazine-5-one (I), 6-methyl-3-thio-1,2,4-triazine-5-one (II), and 2,4-dimethoxy-6-methyl-1,3,5-triazine (III) in dimethylformamide was investigated. Electrochemical techniques including differential pulse voltammetry (DPV), cyclic voltammetry (CV), chronoamperometry, and coulometry were employed to study the mechanism of the electrode process. From the analysis of the voltammetric and spectroscopic experiments a mechanism was proposed for the electroreduction of thio-triazine and triazine compounds. Compounds I and II having thiol groups exhibited similar redox behavior. Both compounds displayed two cathodic peaks, whereas the third compound (III), with no thiol group, showed only one cathodic peak in the same potential range as the second peak of compounds I and II. The results of this study show that in the former wave, the one electron reduction of thiol led to a dimer (disulfide) species and in the latter, the triazine ring was reduced in a two-electron process. The effects of various physical and electrochemical parameters were studied and the electrochemical behavior of the monomers was reported as a function of these parameters. A completely irreversible behavior was observed from cyclic voltammograms obtained under different conditions. Furthermore, in this study some numerical constants, such as diffusion constant, transfer coefficient, and rate constant of coupled chemical reaction were determined. 2005 Elsevier B.V. All rights reserved. Keywords: Triazine derivatives; Glassy carbon electrode; Non-aqueous media; Differential pulse voltammetry; Cyclic voltammetry; Chronoamperometry; Coulometry
1. Introduction The growing interest in the use of pesticides in agricultural production and the negative environmental impacts associated with these compounds has led to the development of different electroanalytical methods [1–5] and electrokinetic studies [6–9]. Additionally, the sensitivity of the electrochemical methods and their applicability over an unusually wide concentration * Corresponding authors. Tel.: +98 021 5901021–51x4612; fax: +98 021 6153397. E-mail addresses: [email protected] (N. Farzinnejad), [email protected] (A.A. Miran Beigi).
0022-0728/$ - see front matter 2005 Elsevier B.V. All rights reserved. doi:10.1016/j.jelechem.2005.03.032
range are highly satisfactory as far as toxicological, ecotoxicological, and environmental regulations are concerned [10]. Triazines are a group of chemically similar herbicides including atrazine, cyanazine, and propazine primarily used to control broadleaf weeds and grasses in the world. All triazine herbicides, including atrazine, cyanazine, and propazine may be released into the environment through effluents discharge from manufacturing facilities and their use as herbicides. They are somewhat persistent in water and mobile in soil. Because of their water solubility, they may leach into the ground water [11,12], and be transported in surface runoff [13].
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The US EPA (United States Environmental Protection Agency) considers atrazine a systemic toxin, which has a potential of causing a variety of acute health effects including congestion of the heart, lungs, and kidneys, hypotension, antidiuresis, muscle spasms, weight loss, and adrenal degeneration. Upon chronic exposure, atrazine also has the potential of causing weight loss, cardiovascular damage, retinal degeneration, and mammary tumors [14]. Similar health effects have been reported with exposure to other triazine compounds [15,16]. Given the increased incidence of mammary gland tumors in female laboratory animals exposed to triazine herbicides, these compounds are classified in group C and are therefore considered possible human carcinogens [17]. Due to the importance of triazine compounds containing thiols and those without thiol groups such as metribuzin and metamitron in toxicological studies, chemical and food industries, and biomedical redox systems, these compounds have been extensively studied for many years [18–24]. However, triazine compounds including 4-amino-6-methyl-3-thio-1,2,4-triazine-5-one and 6-methyl-3-thio-1,2,4-triazine-5-one have rarely been studied especially at glassy carbon (Gc) electrode and in dimethylformamide (DMF) Medium. In the previous investigations, some herbicides containing triazine ring were studied in soils [25–27], plants [28–30], and water [31–36] by voltammetry using, mercury as working electrode in most cases. The electrode reaction of 1,3,5-triazine compounds depends primarily on the species located at the 2-position of the triazine ring, which is a thiomethyl group for some triazines. Therefore, other substituents linked to the ring in different triazines are too far from the electroactive center (2-position) and no pronounced effect on the peak potentials is expected. In this paper, electroreduction of some triazine compounds, namely 4-amino-6-methyl-3-thio-1,2,4-triazine5-one (I), 6-methyl-3-thio-1,2,4-triazine-5-one, (II) and 2,4-dimethoxy-6-methyl-1,3,5-triazine (III) was investigated (as shown in Scheme 1). A variety of electrochemical methods, including differential pulse voltammetry (DPV), cyclic voltammetry (CV), chronoamperometry, (CA) and coulometry were used in DMF at Gc. The reduction mechanism and some numerical parameters, such as diffusion coefficients, transfer coefficients, and rate constant of coupled chemical reaction are also reported.
2. Experimental 2.1. Reagents and materials All reagents and solvents obtained from Merck (Darmstadt, Germany) and Fluka (Switzerland) were
NH2 HS
H HS
N4
O
N1
CH3
O
N1
CH3
2
2
N
N4 N
Test compound I
Test compound II
4-amino-6-methyl-3-thio-1, 2, 4-triazine-5-one
6-methyl-3-thio-1, 2, 4-triazine-5-one
3
N
MeO 5
N
OMe N1
CH 3 Test compound III 2, 4-dimethoxy-6-methyl-1, 3, 5-triazine
Scheme 1. The structure of test compounds I, II, and III.
of analytical grade and were used without further purification. Tetrabuthylammonium perchlorate (TBAP) and N,N-dimethylformamide, DMF, were used as supporting electrolyte and working medium, respectively. In most cases, the electrochemical experiments were performed in 0.10 M TBAP containing 0.001–0.008 M of the triazine compound at 298 K unless otherwise indicated. Argon (Roham Gas Co.) was used as the purging gas after passage through a moisture and organic matter absorbent (conc. H2SO4, Fluka) before introducing to the polarographic cell. All compounds shown in Scheme 1 were synthesized and characterized according to published procedures [37]. Triazines are mostly very toxic and dangerous heteroaromatic compounds [11,16]. They have been recognized as human carcinogens [14,17] and can cause acute health effects on prolonged exposure [15,38]. 2.2. Apparatus A Metrohm model 746 VA Trace Analyzer connected to a 747 VA Stand was used in all experiments. A double wall three-electrode cell (100 ml capacity), which could be thermostated at 298 ± 0.1 K, was used for the voltammetric studies. The Ag/AgCl reference electrode was kept in 3 M KCl saturated with silver ions, supplied by Metrohm Company. Working electrode was a glassy carbon (Gc) electrode (o.d.: 2 mm) press-fitted in a Teflon tube. A platinum wire was also applied as the counter electrode. Coulometric measurements were performed using a ZAG CHEMIE model BHP 2000 microprocessor potentiostat/galvanostat equipped with a feedback circuit to compensate the ir drop. A double beam UV–vis spectrophotometer Unicam model 8700 series was used to identify the disulfide species formed on the surface of the electrode.
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2.3. Procedure The working Gc electrode was cleaned by alternatively polishing with emery paper (10 micron) or graded alumina powder with the same particle size for 5 min before use. It was also rinsed with twice distilled water and subsequently scanned eight times in the range of 1.20 to 2.80 V in 0.10 mol L1 NaHCO3. In order to prevent electrocatalyzed O2 reduction at the negative potential region, special care must be taken in deoxygenation of the test solutions by bubbling purified argon through. In DPV, the test solutions introduced to the cell contained different concentrations of the investigated compounds (1 · 103 to 8 · 103 mol L1). The optimal scan rate and pulse height were 20 mV s1 and 6 mV, respectively. Although the peak currents increase in greater pulse heights and scan rates, but the width of the peaks also increase, which causes a decrease in the selectivity or resolution of cathodic peaks. In CV, high scan rates up to 800 mV s1 were used in survey on the possibility of the presence of competitive electrochemical reactions. Furthermore, repetitive sweeps were carried out in the potential range and in various concentrations of the test compounds. Before running and recording each voltammogram, the electrode surface was cleaned. Double-potential step chronoamperometry measurements were conducted using 4 · 103 mol L1 of each compound in 0.10 mol L1 TBAP in DMF at Gc. The initial potential (Ei), where no electrolysis occurs and the step potential (Es) after an internal time, s = 4 or 5 s, were obtained from each cyclic voltammogram. The number of electrons transferred in the cathodic peaks was obtained by constant potential coulometry. The electrolyte used in these studies was a 30 ml solution of DMF containing 0.10 mol L1 TBAP. The reference solution, containing 0.10 mol L1 TBAP in DMF, was initially electrolyzed and the electrolysis was repeated after addition of the test compound (nearly 2 · 104 mol L1) to the same solution. CV voltammograms obtained in the first and at the end of electrolysis were used to estimate the number of electrons consumed in the cathodic process.
3. Results and discussion 3.1. Differential pulse voltammetry (DPV) Fig. 1 shows typical voltammograms of sample blank and the test compounds I, II, and III in different concentrations in DMF. The sample bank contained 0.10 M TBAP in DMF. In the reverse potentiodynamic scan, we did not find any oxidation wave at Gc in the potential range of 1.20 to 2.80 V, but well-defined irreversible waves were observed in a cathodic scan at potentials of
247
1.65 and 2.40 V, respectively. The test compounds I and II each showed two waves, but in the case of test compound III (with no thiol group) only one wave was observed at nearly 2.40 V. Therefore, considering the results obtained by comparison of the voltammograms and the chemical structure of investigated compounds indicates that half-wave potentials 1.65 and 2.40 V (vs. Ag/AgCl) was, respectively, due to reduction of thiol group and electrochemical cleavage of p-bonding in triazine ring. There was a linear relationship between concentration of the compounds and the corresponding peak currents. In other words, the electrode reaction at Gc was a completely diffusion controlled process. A study on the effect of pulse height and scan rate was carried out. In pulse heights and scan rates greater than the optimal level, an increase in both peak current and peak width was observed which caused a decrease in the selectivity. The scan rate (m) and step time (s, time of sampling currents) are inversely proportional. If s becomes very small, i.e. less than 1 ms, the capacitance current caused by the pulse applied will not decay completely. Since the peak potentials (Ep) in the DPV are functions of the total ionic strength of the solution, Ep was shifted to a more negative value on increasing the concentration of the supporting electrolyte and the test compounds. For example, it shifted approximately 0.10 V in the negative direction for each 8-fold increase in test compound concentration. This was also due to irreversibility of the voltammograms and increase in the viscosity of solution or change in the size of the solvated species. Furthermore, by increasing cell temperature, the peak currents increased approximately 2.0–3.5% for each degree centigrade. 3.2. Cyclic voltammetry (CV) As shown in Fig. 2, typical cyclic voltammograms of the described compounds were obtained in DMF containing 0.10 mol L1 TBAP at Gc electrode. For compounds I and II, two irreversible cathodic waves appeared in the given potential limits while compound III showed only a single wave which irreversibly appeared at nearly 2.50 V. The peak potentials of all three compounds are summarized in Table 1. Therefore, the results obtained by comparison of the voltammograms and the chemical structure of investigated compounds indicate that half-wave potential 1.65 and 2.45 V (vs. Ag/AgCl) were due to reduction of thiol group and electrochemical cleavage of p-bonding in triazine ring, respectively. As expected from DPV studies, a linear increase in the peak heights proportional to increase in concentrations of the compounds at milimolar levels (i.e. 1 · 103 to 8 · 103 mol L1) was observed. Similarly, the electron transfer processes were also controlled via diffusion phenomena. It should be noted that the cell temperature was completely
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Fig. 1. Typical voltammogram using DPV for: (a) sample blank (0.10 M TBAP in DMF); (b) compound I (c1 = 0.003 M, c2 = 0.004 M, c3 = 0.005 M); compound II (c1 = 0.002 M, c2 = 0.004 M, c3 = 0.006 M, c4 = 0.008 M); compound III (c1 = 0.001 M, c2 = 0.002 M, c3 = 0.003 M, c4 = 0.004 M); Ei = 1.20 V; pulse height, 6.0 mV; scan rate, 20 mV s1; purge time, 5 min; cathodic scan.
adjusted at 298 ± 0.10 K to avoid any change in the waveforms. Electrochemical behavior of thiol compounds has been well documented [20,24]. A literature survey on electrochemical reduction of thiol herbicides suggests a mechanism based on one electron reduction of thiol group followed by dimerization reaction to form a disulfide (EC mechanism). Compounds I and II each having a thiol group exhibited similar redox behavior: RC ¼ S þ e þ Hþ ! RCHS0 2RCHS0 ! RCHSSCHR According to the above-proposed reaction scheme, the process involves the uptake of one electron or one H+ ion per herbicide molecule (two electrons or two protons per dimer molecule). The principal products obtained via the radicals produced in the charge transfer step are soluble dimers. Comparison of the voltammograms obtained for compounds I and II (compounds containing a thiol group) and compound III (without a thiol group) shows that the first cathodic peak at 1.65 V (observed only
in DPV or CV of compounds I and II) is related to the reduction of thiol as mentioned above. Therefore, it can be concluded that the first cathodic peak can follow the above EC mechanism, which is well defined in the reviewed literature. On the other hand, there are a few reports [18,19,21–23] on the electrochemical reduction of 1,3,5-triazine derivatives involving reducible azomethine (C@N–) bonds. For example, Ludvik et al. [18] studied electrochemical reduction of metamitron (4-amino-3methyl-6-phenyl-1,2,4-triazin-5(4H)-one) on static mercury drop electrode (SMDE) in DMF and several other solvents. They explained that the herbicide was reduced electrochemically in two 2e steps. Reduction in both steps involved the protonated from of azomethine bonds in 1,6- and 2,3-positions. In addition, similar reports [19,21,22] in relation with 2e-reduction of azomethine in triazine ring have been published. These processes involve the uptake of four electrons and one or two H+ ions per herbicide molecule yielding non-aromatic products. Here, the first cathodic peak appeared only in the DPV or CV of both compounds I and II, whereas the second cathodic peak appeared in the voltammograms of each of compounds. We assumed that the first peak
N. Farzinnejad et al. / Journal of Electroanalytical Chemistry 580 (2005) 245–254
249
(b)
(b)
(a)
(a)
(b) (a)
Fig. 2. Comparison between CV voltammograms of the investigated compounds: (a) sample blank (0.10 M TBAP in DMF); (b) compound I (c = 0.005 M); compound II (c = 0.004 M); compound III (c = 0.001 M); scan rate, 100 mV s1; purge time, 5 min; cathodic scan.
Table 1 The peak potentials of all three compounds Test compound
Ep (V) [DP method]
Ep (V) [CV method]
I (First peak) I (Second peak) II (First peak) II (Second peak) III
1.61 2.40 1.70 2.37 2.43
1.70 2.50 1.62 2.43 2.48
may correspond to the formation of disulfide in an EC mechanism, and the second peak is due to the reduction of azomethine bond. The second cathodic peak appeared in the voltammograms of compounds I, II, and III. The reduction takes place by a two-electron and two-proton transfers. Thus, we propose that the reducible group is the N(1)@C(6) double bond (Scheme 1). This is consistent with the following: the negative charge on N(1) is lower than that on N(2) in test compounds I and II. Furthermore, the negative charge on N(1) is lower than that on N(3) and N(5) in test compound III. This means that is more difficult to reduce the double bond in test compounds I and II N(2)@C(3) and in test compound III N(3)@C(2) and N(5)@C(4). Thus, it is clear why in all three compounds the reducible group is the N(1)@C(6) double bond.
In order to ensure from chemical reaction along with electron transfer process (EC mechanism), the current function (jIp/m1/2j) versus scan rate was plotted for both the first and second cathodic peaks in the investigated compounds. The peak currents (Ip) were obtained from CV studies at various scan rates for each compound. By increasing scan rate, the peak potentials (Ep) related to both peaks were shifted to more negative values due to the irreversibility of charge transfer step in the electrode process. Greater shifts (i.e. DEp > 50 mV) in Ep were observed at scan rates higher than 600 mV s1, which indicates a more irreversible electrochemical process at the surface of the Gc electrode. As seen in Fig. 3(a), a significant decrease in current function of the first peak was observed, by increasing scan rate which confirms the presence of electron transfer process followed by a chemical reaction (EC mechanism), while the current function of the second peak (Fig. 3(b)) was constant throughout the range of the scan rate (50–600 mV s1). This explains an irreversible electron transfer with no coupled chemical reaction (E mechanism) for the more negative peak. The behaviors of the compounds I and II were similar, but the compound III exhibited only E mechanism. As Andrieux and Saveant [39] described, the transfer coefficients (a) for an overall irreversible electron
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(First peak)
(a)
(First peak)
(a)
-1.85
1.1 1.05
0.95
Ep (V)
1/2
Ip /υ
y = -0.0718x - 1.5959 R2 = 0.99
-1.8
1
0.9
-1.75
0.85
-1.7
0.8 0.75
-1.65
0.7 0
100
200
300
400
500
1.6
600
1.8
2
(Second peak)
2.4
2.6
2.8
(Second peak)
(b)
3
-2.5
2.5
-2.49
Ep (V)
I p / υ1/2
(b)
2.2
Log υ
υ
2
y = -0.0422x - 2.3768 R2 = 0.99
-2.48 -2.47
1.5
-2.46 1 50
150
250
350 υ
450
550
-2.45
650
1.9
2.1
2.3
2.5
2.7
Log υ
Fig. 3. Plot of the current function (Ip/m1/2) versus scan rate for first and second cathodic peaks of compound I under the optimal conditions.
Fig. 4. Plot of Ep vs. log m for first and second peaks of compound I, under the same conditions.
transfer can be calculated according to the following equation:
data indicate the involvement of two electrons in the rate-determining step.
Ep ¼ ðb=2Þ log m þ constant;
3.3. Double-potential step chronoamperometry
ð1Þ
where b is the Tafel slope and the intercept of the plot of Ep vs. log m is constant. The slope of the linear regression is equal to b/2 = 0.059/an. Thus, b is 2(0.059/an). The plots of Ep vs. log m are shown in Figs. 4(a) and (b) for the first and second peaks of compound I, respectively. The transfer coefficients can be also obtained according to Eq. (2) [40]: a ¼ 1.857RT =nF ðEp Ep=2 Þ;
ð2Þ
where Ep/2 is the half-peak potential and can preferably be calculated from DPV studies. The transfer coefficients can be also obtained by plots of log Ip vs. log m. The slope of the linear regression is equal to a. The calculated transfer coefficients using the linear plots (Ep vs. log m) and (log Ip vs. log m) are compared with the results obtained from Eq. (2) for each of compounds I, II, and III shown in Table 2. The results in Table 2 indicate that a reasonable value for a in the neighborhood of 0.5. The
Since chronoamperometry is one of the most powerful techniques in investigating EC mechanism, it was used in the course of our studies. Some numerical constants such as diffusion coefficient and rate constant of coupled chemical reaction in the first reduction peak were also determined using this technique. In the absence of a coupled chemical reaction, the cathodic current, ic, is equal with anodic current, ia, considering step (Es) and final (Ef) potentials. Figs. 5(a) and (b) show double-potential step chronoamperograms of the first cathodic peak of compound I and unique peak of compound III, respectively. Ei, Es and Ef of the related peaks were, respectively, selected at 1.0, 1.7 and 1.0 V (for compound I) and at 2.0, 2.5 and 2.0 V (for compound III). As observed, a decrease in ia/ic ratio with increasing time appeared in the chronoamperograms of both compounds I and II at the first cathodic peak. This confirms the
N. Farzinnejad et al. / Journal of Electroanalytical Chemistry 580 (2005) 245–254
251
Table 2 Comparison between the transfer coefficients obtained from Eq. (1), Eq. (2) and the slope of log Ip vs. log t for the available cathodic peaks of compounds I, II, and III Transfer coefficient value, a Compound I
Using Eq. (1) Using Eq. (2) log Ip vs. log t
Compound II
First peak
Second peak
First peak
Second peak
Unique peak
0.41 0.39 0.37
0.35 0.30 0.47
0.44 0.56 0.43
0.35 0.30 0.49
0.34 0.59 0.50
consumption of the product of electron transfer reaction in the following chemical reaction over a long period of time (governed by EC mechanism). The ratio of ia/ic for the second cathodic peak of compounds I and II and the unique peak of compound III were equal to unity. This is consistent with an electron (a) -40 -30 -20
Ip ( A )
i a /ic = ib /i f < 1 -10 0 10 20
transfer process without the presence of any coupled chemical reaction (governed by E mechanism). The results obtained from Chronoamperometry of the investigated herbicides verified the early studies performed in this work. The rate constant of chemical reaction, k, can be obtained by using the working curves, which show ia/ic as a function of ks and t s/s [41]. The rate constants of the formation of disulfide after one electron transfer in the first cathodic peak of compounds I and II were about 0.14 and 0.07 s1, respectively, using the working curves. On the other hand, it was found that difference between the anodic and cathodic currents (Di = ia ic) in compound I was appreciably greater than that in compound II. Apparently, higher constant rates of the disulfide formation are the reason for Di value increases. The diffusion coefficients can also be obtained by chronoamperometry. The relation between current decay and time in a solution of known concentration is given by Cottrel equation: I p ¼ nFAD1=2 C 0 =p1=2 t1=2 ¼ Kt1=2 .
30 0
2
4
6
8
10
t (sec) (b) -50 -40 -30 -20
I p ( A)
Compound III
ia /ic = ib /if = 1
-10 0 10 20 30 40 50 0
2
4
6
8
ð3Þ
1/2
The plots of Ipt vs. t for known concentration of each test compound I, II, and III explain that Ipt1/2 function is constant and time-independent over a wide length of time, especially at long times. The diffusion coefficients obtained from these plots were (1.23 · 105, 0.65 · 105 cm2 s1), (0.77 · 105, 0.69 · 105 cm2 s1), and 0.33 · 106 cm2 s1 for the studied compounds I, II, and III, respectively. In addition, the plots showed that the electron transfer processes for all compounds were diffusion-controlled. In general, the diffusion currents of most substances are relatively smaller in nonaqueous solvents than in aqueous media. Decreased diffusion coefficients may result from either increased viscosity or change in the size of the solvated species. Since in different cases these two effects many predation of the separate influence of each is not a simple matter. 3.4. Coulometry
t (sec) Fig. 5. The double-potential step chronoamperograms of (a) first cathodic peak of compound I; Ei = 1.0 V, Es = 1.7 V, and Ef = 1.0 V, s = 5 s; (b) unique peak of compound III; Ei = 2.0 V, Es = 2.5 V, and Ef = 2.0 V, s = 4 s.
Coulometric experiments were carried out to determine the number of electrons transferred in the reduction process. The potential controlled was selected at 1.70 V vs. Ag/AgCl; much more negative than the first
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peaks displayed in the voltammetric curves of compounds I and II. Under this condition, the electrolysis proceeded completely and an overall number of 1.0 ± 0.1 mol of electron per mol of compound was obtained for both compounds I and II. Furthermore, the overall number of moles of electrons obtained on the second cathodic peak at 2.6 V was 3.0 ± 0.1 per mol of compounds I and II. Considering one electron in the first step, the second peaks involved a 2e reduction process. Coulometry at the controlled-potential of 2.55 V in compound III, showed the an overall electron transfer of 2.0 ± 0.1 electron per molecule which was in good agreement with the second peak in compounds I and II.
3
(a) 2.5
(b) 2
Absorbance
252
1.5
1
0.5
0
3.5. UV–vis spectrophotometry
250
270
290
310
330
350
nm
UV–visible spectra of electrolytic reduction products can help to explain disulfide bond formation resulting from increases in p-bonding electrons and the conjugated system. UV–vis spectrophotometry was used to provide more detailed information on the elucidation of the electrode process. Before employing coulometry in a potential controlled (i.e. 1.70 V), the compounds I and II showed well-defined absorption maxima at 270 and 273 nm, respectively, without any band in the range 300–400 nm. A review of the related literature indicated [20,42], that the compounds containing –S–S– bond show a long-tail absorption band at nearly 310 nm. When using coulometry at 1.70 V, a bath chromic effect at the maximum wavelength was observed and the absorption peak at 270 nm gradually decreased with progress in electrolysis and its location shifted toward longer wavelength near 330 nm. This is a good indication of the formation of disulfide during the electrolysis at a potential more negative than the first cathodic peak of compounds I and II. The UV–vis spectra of compound I (nearly 2 · 104 mol L1) is shown in Fig. 6 before and after coulometry at 1.70 V. Compound II showed similar behavior. The UV spectra of solutions of compound III did not show any absorption peak at 330 nm before and after electrolysis due to the absence of thiol group.
Fig. 6. Absorption spectroscopic characterization of compound I: (a) before electrolysis at E = 1.70 V, kmax = 270 nm; (b) after electrolysis, kmax = 330 nm.
According to the proposed reaction pathway (Fig. 7), thiol compounds I and II show one electron reduction waves due to the formation of dimer (B). There is a good agreement between theoretical and experimental n values obtained from coulometry (1 ± 0.1). Furthermore, spectrophotometric studies confirmed the formation of disulfide in the first reduction. It was proved that the appearance of a long tail absorption peak at 330 nm was related to the –S–S– bond [42,43]. Formation of B would consume one mol of electron per mol of A. At the second reduction wave of compounds I and II, the azomethine in each monomer was reduced in a 2e
H N
HS N
H O
.S + e- + H+
N
CH 3
H
H N
N
O
N
CH 3
A
H O
N
H 3C
N
H S
S
H
H
N
O
N
CH 3
2A
3.6. Mechanism On the basis of the above results, thio-triazine compounds I and II are reduced in a two potential region. As previously described, the current function in the first cathodic peak decreased with increasing scan rate. Furthermore, the ratio of ia/ic in the chronoamperometry of the first peak was less than one, which indicates an EC mechanism. Therefore, it is concluded that the reduction in the first cathodic peak followed the well-defined EC mechanism in thiol herbicides (as shown in Fig. 7).
N
N
B
H O B +4e - + 4H+
H H 3C
N N N H
H S
S
H
H
N N
N
O H CH 3
H
Fig. 7. Proposed reaction pathway for the reduction of compound II on the Gc electrode in DMF Media.
N. Farzinnejad et al. / Journal of Electroanalytical Chemistry 580 (2005) 245–254
process (four electrons per mol of dimer, or two electrons per mol of monomer), whereas the azomethine in compound III was reduced in only one reduction step. The fact that the current function, Ip/m1/2, decreases as the scan rate, t, increases indicates a chemical reaction which is coupled to the electrode process. This is consistent with the proposed mechanism (EC mechanism). However, the ratio of Ip/m1/2 to m for the second peak of compounds I and II and the unique peak of compound III is constant, which predicts an E mechanism for this reduction step. E mechanism in mentioned peaks was also confirmed by chronoamperometry (the ratio ia/ic = 1). Also, it was observed that in compounds I and II, the peak height of the second peak was two times bigger than the first peak.
4. Conclusion The electroreduction of compounds I and II follows a well-defined mechanism in thio-triazine compounds. At the potential of the first peak, the uptake of one electron gives a radical, which dimerizes to disulfide. The azomethine bond in triazine ring reduces in a 2e process. Compound III shows unique reduction peak due to an azomethine moiety. The diffusion constants (D) and transfer coefficients (a) were obtained for the first and second peak of all three compounds. Furthermore, the rate constants of coupled chemical reaction in EC mechanism were determined.
5. Further work Further studies will be focused on construction of a novel chemical sensor for determination of some triazine compounds used in pesticides and food industries. The obtained results also will be compared with those from several advanced voltammetry techniques.
Acknowledgments The authors thank N. Gedayloo and M. Cheraghali from Alzahra University for their valued support during the completion of this work. Authors gratefully wish to thank Mr. Mehdizadeh from Research Institute of Petroleum Industry for revision of this paper.
References [1] J.H. Mendez, R.C. Martinez, E.R. Gonzalo, J. Electroanal. Chem. 244 (1988) 221. [2] C. Li, B.D. James, R.J. Magee, Electroanalysis 2 (1990) 63.
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[3] J.H. Mendez, R.C. Martinez, E.R. Gonzalo, Electroanalysis 2 (1990) 389. [4] R.C. Martinez, F. Becerra, J.H. Mendez, P.G. Martin, Electroanalysis 2 (1990) 567. [5] M.R. Smyth, J.G. Osteryoung, Anal. Chim. Acta. 96 (1978) 335. [6] P.C. Nascimento, D. Bohrer, L.M. Carvalho, J. Trevisan, J. Braz. Chem. Soc. 14 (2003) 577. [7] M. Chicharro, A. Zapardiel, E. Bermejo, M. Moreno, Talanta 59 (2003) 37. [8] R. Cho, Shokuhin Eiseigaku Zasshi 40 (1999) 396. [9] K. Wechsler, M. Rombourg, A. Exingerand, C. Breuzin, Ann. Fals. Exp. Chim. 924 (1994) 23. [10] J. Barek, A.G. Fogg, J.C. Moreira, M.V.B. Zanoni, J. Zima, Anal. Chim. Acta 320 (1996) 31. [11] E.M. Thurman, D.A. Goolsby, M.T. Meyer, D.W. Kolpin, Environ. Sci. Technol. 25 (1991) 1794. [12] E.M. Thurman, D.A. Goolsby, M.T. Meyer, M.S. Mills, M.L. Pomes, D.W. Kolpin, Environ. Sci. Technol. 26 (1992) 2440. [13] R.A. Grover, Environmental Chemistry of Herbicides, CRC Press, Boca Raton, FL, 1988, pp. 45–88. [14] US EPA, Drinking water and health: technical fact sheet on atrazine, United States Environmental Protection Agency, Washington, DC, 1999. Available from:
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[36] J. Ludvik, P. Zuman, J. Microchem. 64 (2000) 15. [37] M.M. Heravi, P. Khosrofar, J. Sci. I. R. Iran 7 (1996) 86. [38] J.P. Carder, K.D. Hoagland, Envirom. Toxicol. Chem. 17 (1998) 1415. [39] C.P. Andrieux, J.M. Saveant, J. Electroanal. Chem. 93 (1978) 163.
[40] A.J. Bard, R. Faulkner, Electrochemical Methods, John Wiley, New York, 1980. [41] W.M. Schwars, I. Shain, J. Phys. Chem. 69 (1965) 30. [42] H.P. Koch, J. Chem. Soc. (1949) 407. [43] W. Stricks, J. Frischmann, R.G. Mueller, J. Electrochem. Soc. 109 (1962) 518.
Electroanalytical Chemistry Journal of Electroanalytical Chemistry 580 (2005) 245–254 www.elsevier.com/locate/jelechem
Electrochemical behavior of some triazine derivatives at glassy carbon electrode in non-aqueous media N. Farzinnejad
a,*
, A.A. Miran Beigi
b,*
, L. Fotouhi a, K. Torkestani b, H.A. Ghadirian
b
a
b
Department of Chemistry, Faculty of Science, Al-Zahra University, P.O. Box 1993891176, Tehran, Iran Standard Laboratory, Physical and Chemical Analyses Group, Research Institute of Petroleum Industry, P.O. Box 18745-4163, Tehran, Iran Received 12 October 2004; received in revised form 4 March 2005; accepted 17 March 2005 Available online 23 May 2005
Abstract Electroreduction of 4-amino-6-methyl-3-thio-1,2,4-triazine-5-one (I), 6-methyl-3-thio-1,2,4-triazine-5-one (II), and 2,4-dimethoxy-6-methyl-1,3,5-triazine (III) in dimethylformamide was investigated. Electrochemical techniques including differential pulse voltammetry (DPV), cyclic voltammetry (CV), chronoamperometry, and coulometry were employed to study the mechanism of the electrode process. From the analysis of the voltammetric and spectroscopic experiments a mechanism was proposed for the electroreduction of thio-triazine and triazine compounds. Compounds I and II having thiol groups exhibited similar redox behavior. Both compounds displayed two cathodic peaks, whereas the third compound (III), with no thiol group, showed only one cathodic peak in the same potential range as the second peak of compounds I and II. The results of this study show that in the former wave, the one electron reduction of thiol led to a dimer (disulfide) species and in the latter, the triazine ring was reduced in a two-electron process. The effects of various physical and electrochemical parameters were studied and the electrochemical behavior of the monomers was reported as a function of these parameters. A completely irreversible behavior was observed from cyclic voltammograms obtained under different conditions. Furthermore, in this study some numerical constants, such as diffusion constant, transfer coefficient, and rate constant of coupled chemical reaction were determined. 2005 Elsevier B.V. All rights reserved. Keywords: Triazine derivatives; Glassy carbon electrode; Non-aqueous media; Differential pulse voltammetry; Cyclic voltammetry; Chronoamperometry; Coulometry
1. Introduction The growing interest in the use of pesticides in agricultural production and the negative environmental impacts associated with these compounds has led to the development of different electroanalytical methods [1–5] and electrokinetic studies [6–9]. Additionally, the sensitivity of the electrochemical methods and their applicability over an unusually wide concentration * Corresponding authors. Tel.: +98 021 5901021–51x4612; fax: +98 021 6153397. E-mail addresses: [email protected] (N. Farzinnejad), [email protected] (A.A. Miran Beigi).
0022-0728/$ - see front matter 2005 Elsevier B.V. All rights reserved. doi:10.1016/j.jelechem.2005.03.032
range are highly satisfactory as far as toxicological, ecotoxicological, and environmental regulations are concerned [10]. Triazines are a group of chemically similar herbicides including atrazine, cyanazine, and propazine primarily used to control broadleaf weeds and grasses in the world. All triazine herbicides, including atrazine, cyanazine, and propazine may be released into the environment through effluents discharge from manufacturing facilities and their use as herbicides. They are somewhat persistent in water and mobile in soil. Because of their water solubility, they may leach into the ground water [11,12], and be transported in surface runoff [13].
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The US EPA (United States Environmental Protection Agency) considers atrazine a systemic toxin, which has a potential of causing a variety of acute health effects including congestion of the heart, lungs, and kidneys, hypotension, antidiuresis, muscle spasms, weight loss, and adrenal degeneration. Upon chronic exposure, atrazine also has the potential of causing weight loss, cardiovascular damage, retinal degeneration, and mammary tumors [14]. Similar health effects have been reported with exposure to other triazine compounds [15,16]. Given the increased incidence of mammary gland tumors in female laboratory animals exposed to triazine herbicides, these compounds are classified in group C and are therefore considered possible human carcinogens [17]. Due to the importance of triazine compounds containing thiols and those without thiol groups such as metribuzin and metamitron in toxicological studies, chemical and food industries, and biomedical redox systems, these compounds have been extensively studied for many years [18–24]. However, triazine compounds including 4-amino-6-methyl-3-thio-1,2,4-triazine-5-one and 6-methyl-3-thio-1,2,4-triazine-5-one have rarely been studied especially at glassy carbon (Gc) electrode and in dimethylformamide (DMF) Medium. In the previous investigations, some herbicides containing triazine ring were studied in soils [25–27], plants [28–30], and water [31–36] by voltammetry using, mercury as working electrode in most cases. The electrode reaction of 1,3,5-triazine compounds depends primarily on the species located at the 2-position of the triazine ring, which is a thiomethyl group for some triazines. Therefore, other substituents linked to the ring in different triazines are too far from the electroactive center (2-position) and no pronounced effect on the peak potentials is expected. In this paper, electroreduction of some triazine compounds, namely 4-amino-6-methyl-3-thio-1,2,4-triazine5-one (I), 6-methyl-3-thio-1,2,4-triazine-5-one, (II) and 2,4-dimethoxy-6-methyl-1,3,5-triazine (III) was investigated (as shown in Scheme 1). A variety of electrochemical methods, including differential pulse voltammetry (DPV), cyclic voltammetry (CV), chronoamperometry, (CA) and coulometry were used in DMF at Gc. The reduction mechanism and some numerical parameters, such as diffusion coefficients, transfer coefficients, and rate constant of coupled chemical reaction are also reported.
2. Experimental 2.1. Reagents and materials All reagents and solvents obtained from Merck (Darmstadt, Germany) and Fluka (Switzerland) were
NH2 HS
H HS
N4
O
N1
CH3
O
N1
CH3
2
2
N
N4 N
Test compound I
Test compound II
4-amino-6-methyl-3-thio-1, 2, 4-triazine-5-one
6-methyl-3-thio-1, 2, 4-triazine-5-one
3
N
MeO 5
N
OMe N1
CH 3 Test compound III 2, 4-dimethoxy-6-methyl-1, 3, 5-triazine
Scheme 1. The structure of test compounds I, II, and III.
of analytical grade and were used without further purification. Tetrabuthylammonium perchlorate (TBAP) and N,N-dimethylformamide, DMF, were used as supporting electrolyte and working medium, respectively. In most cases, the electrochemical experiments were performed in 0.10 M TBAP containing 0.001–0.008 M of the triazine compound at 298 K unless otherwise indicated. Argon (Roham Gas Co.) was used as the purging gas after passage through a moisture and organic matter absorbent (conc. H2SO4, Fluka) before introducing to the polarographic cell. All compounds shown in Scheme 1 were synthesized and characterized according to published procedures [37]. Triazines are mostly very toxic and dangerous heteroaromatic compounds [11,16]. They have been recognized as human carcinogens [14,17] and can cause acute health effects on prolonged exposure [15,38]. 2.2. Apparatus A Metrohm model 746 VA Trace Analyzer connected to a 747 VA Stand was used in all experiments. A double wall three-electrode cell (100 ml capacity), which could be thermostated at 298 ± 0.1 K, was used for the voltammetric studies. The Ag/AgCl reference electrode was kept in 3 M KCl saturated with silver ions, supplied by Metrohm Company. Working electrode was a glassy carbon (Gc) electrode (o.d.: 2 mm) press-fitted in a Teflon tube. A platinum wire was also applied as the counter electrode. Coulometric measurements were performed using a ZAG CHEMIE model BHP 2000 microprocessor potentiostat/galvanostat equipped with a feedback circuit to compensate the ir drop. A double beam UV–vis spectrophotometer Unicam model 8700 series was used to identify the disulfide species formed on the surface of the electrode.
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2.3. Procedure The working Gc electrode was cleaned by alternatively polishing with emery paper (10 micron) or graded alumina powder with the same particle size for 5 min before use. It was also rinsed with twice distilled water and subsequently scanned eight times in the range of 1.20 to 2.80 V in 0.10 mol L1 NaHCO3. In order to prevent electrocatalyzed O2 reduction at the negative potential region, special care must be taken in deoxygenation of the test solutions by bubbling purified argon through. In DPV, the test solutions introduced to the cell contained different concentrations of the investigated compounds (1 · 103 to 8 · 103 mol L1). The optimal scan rate and pulse height were 20 mV s1 and 6 mV, respectively. Although the peak currents increase in greater pulse heights and scan rates, but the width of the peaks also increase, which causes a decrease in the selectivity or resolution of cathodic peaks. In CV, high scan rates up to 800 mV s1 were used in survey on the possibility of the presence of competitive electrochemical reactions. Furthermore, repetitive sweeps were carried out in the potential range and in various concentrations of the test compounds. Before running and recording each voltammogram, the electrode surface was cleaned. Double-potential step chronoamperometry measurements were conducted using 4 · 103 mol L1 of each compound in 0.10 mol L1 TBAP in DMF at Gc. The initial potential (Ei), where no electrolysis occurs and the step potential (Es) after an internal time, s = 4 or 5 s, were obtained from each cyclic voltammogram. The number of electrons transferred in the cathodic peaks was obtained by constant potential coulometry. The electrolyte used in these studies was a 30 ml solution of DMF containing 0.10 mol L1 TBAP. The reference solution, containing 0.10 mol L1 TBAP in DMF, was initially electrolyzed and the electrolysis was repeated after addition of the test compound (nearly 2 · 104 mol L1) to the same solution. CV voltammograms obtained in the first and at the end of electrolysis were used to estimate the number of electrons consumed in the cathodic process.
3. Results and discussion 3.1. Differential pulse voltammetry (DPV) Fig. 1 shows typical voltammograms of sample blank and the test compounds I, II, and III in different concentrations in DMF. The sample bank contained 0.10 M TBAP in DMF. In the reverse potentiodynamic scan, we did not find any oxidation wave at Gc in the potential range of 1.20 to 2.80 V, but well-defined irreversible waves were observed in a cathodic scan at potentials of
247
1.65 and 2.40 V, respectively. The test compounds I and II each showed two waves, but in the case of test compound III (with no thiol group) only one wave was observed at nearly 2.40 V. Therefore, considering the results obtained by comparison of the voltammograms and the chemical structure of investigated compounds indicates that half-wave potentials 1.65 and 2.40 V (vs. Ag/AgCl) was, respectively, due to reduction of thiol group and electrochemical cleavage of p-bonding in triazine ring. There was a linear relationship between concentration of the compounds and the corresponding peak currents. In other words, the electrode reaction at Gc was a completely diffusion controlled process. A study on the effect of pulse height and scan rate was carried out. In pulse heights and scan rates greater than the optimal level, an increase in both peak current and peak width was observed which caused a decrease in the selectivity. The scan rate (m) and step time (s, time of sampling currents) are inversely proportional. If s becomes very small, i.e. less than 1 ms, the capacitance current caused by the pulse applied will not decay completely. Since the peak potentials (Ep) in the DPV are functions of the total ionic strength of the solution, Ep was shifted to a more negative value on increasing the concentration of the supporting electrolyte and the test compounds. For example, it shifted approximately 0.10 V in the negative direction for each 8-fold increase in test compound concentration. This was also due to irreversibility of the voltammograms and increase in the viscosity of solution or change in the size of the solvated species. Furthermore, by increasing cell temperature, the peak currents increased approximately 2.0–3.5% for each degree centigrade. 3.2. Cyclic voltammetry (CV) As shown in Fig. 2, typical cyclic voltammograms of the described compounds were obtained in DMF containing 0.10 mol L1 TBAP at Gc electrode. For compounds I and II, two irreversible cathodic waves appeared in the given potential limits while compound III showed only a single wave which irreversibly appeared at nearly 2.50 V. The peak potentials of all three compounds are summarized in Table 1. Therefore, the results obtained by comparison of the voltammograms and the chemical structure of investigated compounds indicate that half-wave potential 1.65 and 2.45 V (vs. Ag/AgCl) were due to reduction of thiol group and electrochemical cleavage of p-bonding in triazine ring, respectively. As expected from DPV studies, a linear increase in the peak heights proportional to increase in concentrations of the compounds at milimolar levels (i.e. 1 · 103 to 8 · 103 mol L1) was observed. Similarly, the electron transfer processes were also controlled via diffusion phenomena. It should be noted that the cell temperature was completely
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Fig. 1. Typical voltammogram using DPV for: (a) sample blank (0.10 M TBAP in DMF); (b) compound I (c1 = 0.003 M, c2 = 0.004 M, c3 = 0.005 M); compound II (c1 = 0.002 M, c2 = 0.004 M, c3 = 0.006 M, c4 = 0.008 M); compound III (c1 = 0.001 M, c2 = 0.002 M, c3 = 0.003 M, c4 = 0.004 M); Ei = 1.20 V; pulse height, 6.0 mV; scan rate, 20 mV s1; purge time, 5 min; cathodic scan.
adjusted at 298 ± 0.10 K to avoid any change in the waveforms. Electrochemical behavior of thiol compounds has been well documented [20,24]. A literature survey on electrochemical reduction of thiol herbicides suggests a mechanism based on one electron reduction of thiol group followed by dimerization reaction to form a disulfide (EC mechanism). Compounds I and II each having a thiol group exhibited similar redox behavior: RC ¼ S þ e þ Hþ ! RCHS0 2RCHS0 ! RCHSSCHR According to the above-proposed reaction scheme, the process involves the uptake of one electron or one H+ ion per herbicide molecule (two electrons or two protons per dimer molecule). The principal products obtained via the radicals produced in the charge transfer step are soluble dimers. Comparison of the voltammograms obtained for compounds I and II (compounds containing a thiol group) and compound III (without a thiol group) shows that the first cathodic peak at 1.65 V (observed only
in DPV or CV of compounds I and II) is related to the reduction of thiol as mentioned above. Therefore, it can be concluded that the first cathodic peak can follow the above EC mechanism, which is well defined in the reviewed literature. On the other hand, there are a few reports [18,19,21–23] on the electrochemical reduction of 1,3,5-triazine derivatives involving reducible azomethine (C@N–) bonds. For example, Ludvik et al. [18] studied electrochemical reduction of metamitron (4-amino-3methyl-6-phenyl-1,2,4-triazin-5(4H)-one) on static mercury drop electrode (SMDE) in DMF and several other solvents. They explained that the herbicide was reduced electrochemically in two 2e steps. Reduction in both steps involved the protonated from of azomethine bonds in 1,6- and 2,3-positions. In addition, similar reports [19,21,22] in relation with 2e-reduction of azomethine in triazine ring have been published. These processes involve the uptake of four electrons and one or two H+ ions per herbicide molecule yielding non-aromatic products. Here, the first cathodic peak appeared only in the DPV or CV of both compounds I and II, whereas the second cathodic peak appeared in the voltammograms of each of compounds. We assumed that the first peak
N. Farzinnejad et al. / Journal of Electroanalytical Chemistry 580 (2005) 245–254
249
(b)
(b)
(a)
(a)
(b) (a)
Fig. 2. Comparison between CV voltammograms of the investigated compounds: (a) sample blank (0.10 M TBAP in DMF); (b) compound I (c = 0.005 M); compound II (c = 0.004 M); compound III (c = 0.001 M); scan rate, 100 mV s1; purge time, 5 min; cathodic scan.
Table 1 The peak potentials of all three compounds Test compound
Ep (V) [DP method]
Ep (V) [CV method]
I (First peak) I (Second peak) II (First peak) II (Second peak) III
1.61 2.40 1.70 2.37 2.43
1.70 2.50 1.62 2.43 2.48
may correspond to the formation of disulfide in an EC mechanism, and the second peak is due to the reduction of azomethine bond. The second cathodic peak appeared in the voltammograms of compounds I, II, and III. The reduction takes place by a two-electron and two-proton transfers. Thus, we propose that the reducible group is the N(1)@C(6) double bond (Scheme 1). This is consistent with the following: the negative charge on N(1) is lower than that on N(2) in test compounds I and II. Furthermore, the negative charge on N(1) is lower than that on N(3) and N(5) in test compound III. This means that is more difficult to reduce the double bond in test compounds I and II N(2)@C(3) and in test compound III N(3)@C(2) and N(5)@C(4). Thus, it is clear why in all three compounds the reducible group is the N(1)@C(6) double bond.
In order to ensure from chemical reaction along with electron transfer process (EC mechanism), the current function (jIp/m1/2j) versus scan rate was plotted for both the first and second cathodic peaks in the investigated compounds. The peak currents (Ip) were obtained from CV studies at various scan rates for each compound. By increasing scan rate, the peak potentials (Ep) related to both peaks were shifted to more negative values due to the irreversibility of charge transfer step in the electrode process. Greater shifts (i.e. DEp > 50 mV) in Ep were observed at scan rates higher than 600 mV s1, which indicates a more irreversible electrochemical process at the surface of the Gc electrode. As seen in Fig. 3(a), a significant decrease in current function of the first peak was observed, by increasing scan rate which confirms the presence of electron transfer process followed by a chemical reaction (EC mechanism), while the current function of the second peak (Fig. 3(b)) was constant throughout the range of the scan rate (50–600 mV s1). This explains an irreversible electron transfer with no coupled chemical reaction (E mechanism) for the more negative peak. The behaviors of the compounds I and II were similar, but the compound III exhibited only E mechanism. As Andrieux and Saveant [39] described, the transfer coefficients (a) for an overall irreversible electron
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(First peak)
(a)
(First peak)
(a)
-1.85
1.1 1.05
0.95
Ep (V)
1/2
Ip /υ
y = -0.0718x - 1.5959 R2 = 0.99
-1.8
1
0.9
-1.75
0.85
-1.7
0.8 0.75
-1.65
0.7 0
100
200
300
400
500
1.6
600
1.8
2
(Second peak)
2.4
2.6
2.8
(Second peak)
(b)
3
-2.5
2.5
-2.49
Ep (V)
I p / υ1/2
(b)
2.2
Log υ
υ
2
y = -0.0422x - 2.3768 R2 = 0.99
-2.48 -2.47
1.5
-2.46 1 50
150
250
350 υ
450
550
-2.45
650
1.9
2.1
2.3
2.5
2.7
Log υ
Fig. 3. Plot of the current function (Ip/m1/2) versus scan rate for first and second cathodic peaks of compound I under the optimal conditions.
Fig. 4. Plot of Ep vs. log m for first and second peaks of compound I, under the same conditions.
transfer can be calculated according to the following equation:
data indicate the involvement of two electrons in the rate-determining step.
Ep ¼ ðb=2Þ log m þ constant;
3.3. Double-potential step chronoamperometry
ð1Þ
where b is the Tafel slope and the intercept of the plot of Ep vs. log m is constant. The slope of the linear regression is equal to b/2 = 0.059/an. Thus, b is 2(0.059/an). The plots of Ep vs. log m are shown in Figs. 4(a) and (b) for the first and second peaks of compound I, respectively. The transfer coefficients can be also obtained according to Eq. (2) [40]: a ¼ 1.857RT =nF ðEp Ep=2 Þ;
ð2Þ
where Ep/2 is the half-peak potential and can preferably be calculated from DPV studies. The transfer coefficients can be also obtained by plots of log Ip vs. log m. The slope of the linear regression is equal to a. The calculated transfer coefficients using the linear plots (Ep vs. log m) and (log Ip vs. log m) are compared with the results obtained from Eq. (2) for each of compounds I, II, and III shown in Table 2. The results in Table 2 indicate that a reasonable value for a in the neighborhood of 0.5. The
Since chronoamperometry is one of the most powerful techniques in investigating EC mechanism, it was used in the course of our studies. Some numerical constants such as diffusion coefficient and rate constant of coupled chemical reaction in the first reduction peak were also determined using this technique. In the absence of a coupled chemical reaction, the cathodic current, ic, is equal with anodic current, ia, considering step (Es) and final (Ef) potentials. Figs. 5(a) and (b) show double-potential step chronoamperograms of the first cathodic peak of compound I and unique peak of compound III, respectively. Ei, Es and Ef of the related peaks were, respectively, selected at 1.0, 1.7 and 1.0 V (for compound I) and at 2.0, 2.5 and 2.0 V (for compound III). As observed, a decrease in ia/ic ratio with increasing time appeared in the chronoamperograms of both compounds I and II at the first cathodic peak. This confirms the
N. Farzinnejad et al. / Journal of Electroanalytical Chemistry 580 (2005) 245–254
251
Table 2 Comparison between the transfer coefficients obtained from Eq. (1), Eq. (2) and the slope of log Ip vs. log t for the available cathodic peaks of compounds I, II, and III Transfer coefficient value, a Compound I
Using Eq. (1) Using Eq. (2) log Ip vs. log t
Compound II
First peak
Second peak
First peak
Second peak
Unique peak
0.41 0.39 0.37
0.35 0.30 0.47
0.44 0.56 0.43
0.35 0.30 0.49
0.34 0.59 0.50
consumption of the product of electron transfer reaction in the following chemical reaction over a long period of time (governed by EC mechanism). The ratio of ia/ic for the second cathodic peak of compounds I and II and the unique peak of compound III were equal to unity. This is consistent with an electron (a) -40 -30 -20
Ip ( A )
i a /ic = ib /i f < 1 -10 0 10 20
transfer process without the presence of any coupled chemical reaction (governed by E mechanism). The results obtained from Chronoamperometry of the investigated herbicides verified the early studies performed in this work. The rate constant of chemical reaction, k, can be obtained by using the working curves, which show ia/ic as a function of ks and t s/s [41]. The rate constants of the formation of disulfide after one electron transfer in the first cathodic peak of compounds I and II were about 0.14 and 0.07 s1, respectively, using the working curves. On the other hand, it was found that difference between the anodic and cathodic currents (Di = ia ic) in compound I was appreciably greater than that in compound II. Apparently, higher constant rates of the disulfide formation are the reason for Di value increases. The diffusion coefficients can also be obtained by chronoamperometry. The relation between current decay and time in a solution of known concentration is given by Cottrel equation: I p ¼ nFAD1=2 C 0 =p1=2 t1=2 ¼ Kt1=2 .
30 0
2
4
6
8
10
t (sec) (b) -50 -40 -30 -20
I p ( A)
Compound III
ia /ic = ib /if = 1
-10 0 10 20 30 40 50 0
2
4
6
8
ð3Þ
1/2
The plots of Ipt vs. t for known concentration of each test compound I, II, and III explain that Ipt1/2 function is constant and time-independent over a wide length of time, especially at long times. The diffusion coefficients obtained from these plots were (1.23 · 105, 0.65 · 105 cm2 s1), (0.77 · 105, 0.69 · 105 cm2 s1), and 0.33 · 106 cm2 s1 for the studied compounds I, II, and III, respectively. In addition, the plots showed that the electron transfer processes for all compounds were diffusion-controlled. In general, the diffusion currents of most substances are relatively smaller in nonaqueous solvents than in aqueous media. Decreased diffusion coefficients may result from either increased viscosity or change in the size of the solvated species. Since in different cases these two effects many predation of the separate influence of each is not a simple matter. 3.4. Coulometry
t (sec) Fig. 5. The double-potential step chronoamperograms of (a) first cathodic peak of compound I; Ei = 1.0 V, Es = 1.7 V, and Ef = 1.0 V, s = 5 s; (b) unique peak of compound III; Ei = 2.0 V, Es = 2.5 V, and Ef = 2.0 V, s = 4 s.
Coulometric experiments were carried out to determine the number of electrons transferred in the reduction process. The potential controlled was selected at 1.70 V vs. Ag/AgCl; much more negative than the first
N. Farzinnejad et al. / Journal of Electroanalytical Chemistry 580 (2005) 245–254
peaks displayed in the voltammetric curves of compounds I and II. Under this condition, the electrolysis proceeded completely and an overall number of 1.0 ± 0.1 mol of electron per mol of compound was obtained for both compounds I and II. Furthermore, the overall number of moles of electrons obtained on the second cathodic peak at 2.6 V was 3.0 ± 0.1 per mol of compounds I and II. Considering one electron in the first step, the second peaks involved a 2e reduction process. Coulometry at the controlled-potential of 2.55 V in compound III, showed the an overall electron transfer of 2.0 ± 0.1 electron per molecule which was in good agreement with the second peak in compounds I and II.
3
(a) 2.5
(b) 2
Absorbance
252
1.5
1
0.5
0
3.5. UV–vis spectrophotometry
250
270
290
310
330
350
nm
UV–visible spectra of electrolytic reduction products can help to explain disulfide bond formation resulting from increases in p-bonding electrons and the conjugated system. UV–vis spectrophotometry was used to provide more detailed information on the elucidation of the electrode process. Before employing coulometry in a potential controlled (i.e. 1.70 V), the compounds I and II showed well-defined absorption maxima at 270 and 273 nm, respectively, without any band in the range 300–400 nm. A review of the related literature indicated [20,42], that the compounds containing –S–S– bond show a long-tail absorption band at nearly 310 nm. When using coulometry at 1.70 V, a bath chromic effect at the maximum wavelength was observed and the absorption peak at 270 nm gradually decreased with progress in electrolysis and its location shifted toward longer wavelength near 330 nm. This is a good indication of the formation of disulfide during the electrolysis at a potential more negative than the first cathodic peak of compounds I and II. The UV–vis spectra of compound I (nearly 2 · 104 mol L1) is shown in Fig. 6 before and after coulometry at 1.70 V. Compound II showed similar behavior. The UV spectra of solutions of compound III did not show any absorption peak at 330 nm before and after electrolysis due to the absence of thiol group.
Fig. 6. Absorption spectroscopic characterization of compound I: (a) before electrolysis at E = 1.70 V, kmax = 270 nm; (b) after electrolysis, kmax = 330 nm.
According to the proposed reaction pathway (Fig. 7), thiol compounds I and II show one electron reduction waves due to the formation of dimer (B). There is a good agreement between theoretical and experimental n values obtained from coulometry (1 ± 0.1). Furthermore, spectrophotometric studies confirmed the formation of disulfide in the first reduction. It was proved that the appearance of a long tail absorption peak at 330 nm was related to the –S–S– bond [42,43]. Formation of B would consume one mol of electron per mol of A. At the second reduction wave of compounds I and II, the azomethine in each monomer was reduced in a 2e
H N
HS N
H O
.S + e- + H+
N
CH 3
H
H N
N
O
N
CH 3
A
H O
N
H 3C
N
H S
S
H
H
N
O
N
CH 3
2A
3.6. Mechanism On the basis of the above results, thio-triazine compounds I and II are reduced in a two potential region. As previously described, the current function in the first cathodic peak decreased with increasing scan rate. Furthermore, the ratio of ia/ic in the chronoamperometry of the first peak was less than one, which indicates an EC mechanism. Therefore, it is concluded that the reduction in the first cathodic peak followed the well-defined EC mechanism in thiol herbicides (as shown in Fig. 7).
N
N
B
H O B +4e - + 4H+
H H 3C
N N N H
H S
S
H
H
N N
N
O H CH 3
H
Fig. 7. Proposed reaction pathway for the reduction of compound II on the Gc electrode in DMF Media.
N. Farzinnejad et al. / Journal of Electroanalytical Chemistry 580 (2005) 245–254
process (four electrons per mol of dimer, or two electrons per mol of monomer), whereas the azomethine in compound III was reduced in only one reduction step. The fact that the current function, Ip/m1/2, decreases as the scan rate, t, increases indicates a chemical reaction which is coupled to the electrode process. This is consistent with the proposed mechanism (EC mechanism). However, the ratio of Ip/m1/2 to m for the second peak of compounds I and II and the unique peak of compound III is constant, which predicts an E mechanism for this reduction step. E mechanism in mentioned peaks was also confirmed by chronoamperometry (the ratio ia/ic = 1). Also, it was observed that in compounds I and II, the peak height of the second peak was two times bigger than the first peak.
4. Conclusion The electroreduction of compounds I and II follows a well-defined mechanism in thio-triazine compounds. At the potential of the first peak, the uptake of one electron gives a radical, which dimerizes to disulfide. The azomethine bond in triazine ring reduces in a 2e process. Compound III shows unique reduction peak due to an azomethine moiety. The diffusion constants (D) and transfer coefficients (a) were obtained for the first and second peak of all three compounds. Furthermore, the rate constants of coupled chemical reaction in EC mechanism were determined.
5. Further work Further studies will be focused on construction of a novel chemical sensor for determination of some triazine compounds used in pesticides and food industries. The obtained results also will be compared with those from several advanced voltammetry techniques.
Acknowledgments The authors thank N. Gedayloo and M. Cheraghali from Alzahra University for their valued support during the completion of this work. Authors gratefully wish to thank Mr. Mehdizadeh from Research Institute of Petroleum Industry for revision of this paper.
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