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Ferrocene sulfonates as electrocatalysts for sulfide detection
Electrochimica Acta 52 (2006) 499–503
Ferrocene sulfonates as electrocatalysts for sulfide detection Nathan S. Lawrence ∗ , Gary J. Tustin, Michael Faulkner 1 , Timothy G.J. Jones Schlumberger Cambridge Research, High Cross, Madingley Road, Cambridge CB3 0EL, UK Received 21 March 2006; received in revised form 16 May 2006; accepted 16 May 2006 Available online 3 July 2006
Abstract The electrochemical characterization of both the mono- and di-substituted forms of ferrocene sulfonate are given. The results show both species produce voltammograms consistent with quasi-irreversible diffusion controlled redox reactions. The FcSO3 − species was found to be easier to oxidize than its Fc(SO3 )2 2− counterpart, due to the electron withdrawing affect of the sulfonate group on the Fe centre. In the presence of sulfide, the voltammetric response of FcSO3 − is shown to be consistent with the occurrence of an electrocatalytic EC reaction. This analytical response was utilized as a means of determining sulfide and was found to be linear over the concentration 0.02–1 mM with a limit of detection of 14 M. © 2006 Elsevier Ltd. All rights reserved. Keywords: Ferrocene sulfonate; Hydrogen sulfide; Electrocatalytic; Mercaptoethanol; Voltammetry
1. Introduction The determination of sulfide species is important to analytical and environmental chemists [1,2]. This is due to the high toxicity of liberated hydrogen sulfide, as it poses a major problem to those who handle and remove sulfide-contaminated products. The high reactivity of sulfide [3] and hence low accumulative potential of the species minimises the risk of everyday exposure to levels that are likely to cause any serious, adverse health effects. However, there are a number of instances in which localised concentrations of the anion may greatly exceed safety limits and pose significant concerns for those required to work within such environments. Although a wide range of protocols are available for the detection of sulfide which encompass techniques, such as chromatography, spectroscopy and solid state devices [2], electrochemical procedures provide some of the more viable options due to their design simplicity, sensitivity and low cost [4–9]. In the following, we examine the possibility of utilizing the soluble ferrocene derivatives ferrocene monosulfonate (FcSO3 − ) and ferrocene disulfonate (Fc(SO3 )2 2− ) as electrochemical mediators for the detection of sulfide. The synthesis
∗
Corresponding author. Tel.: +44 1223 325224. E-mail address: [email protected] (N.S. Lawrence). 1 Present address: Castle Sixth Form Centre, Kenilworth School, Rouncil Lane, Kenilworth CV8 1FN, UK. 0013-4686/$ – see front matter © 2006 Elsevier Ltd. All rights reserved. doi:10.1016/j.electacta.2006.05.031
of both the mono- and di-substituted ferrocene sulfonates has been described previously [10,11]. Ferrocene sulfonates have been used extensively as anion dopants in the formation of various cationic polymers (for example, polythiophene, polypyrrole [12] and polyaniline [13–15]) or as mediators in enzymatic reactions [16,17]. It can be envisaged that these soluble ferrocene compounds can undergo an electrocatalytic reaction with sulfide and thiol species in an analogous manner to that of the ferrocyanide/ferricyanide redox couple [18–22] and solid ferrocene [23]. In this case, the ferrocene is first oxidized to the ferricenium ion, which is subsequently homogeneously reduced back to ferrocene by the thiol species. This regenerated ferrocence can then be again oxidized at the electrode surface to produce the analytical signal. Herein, we discuss the electrochemistry of various ferrocene sulfonate salts and their subsequent response to hydrogen sulfide. 2. Experimental 2.1. Reagents All chemicals were supplied by Aldrich and used without further purification. The buffered solutions were prepared as follows: pH 4 potassium phthalate buffer (0.05 M), pH 7 disodium hydrogen phosphate (0.025 M) and potassium dihy-
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drogen phosphate (0.025 M) and pH 9 disodium tetraborate (0.05 M). 2.2. Apparatus Electrochemical measurements were recorded using an PGSTAT30 potentiostat (Ecochemie, Netherlands) with a standard three-electrode configuration. A platinum wire (1 mm diameter, Goodfellow Metals, Cambridge, UK) provided the counter electrode and a saturated calomel electrode (SCE, Radiometer, Copenhagen) acted as the reference. The working electrode was composed of either glassy carbon (GC, 2 mm diameter) or boron-doped diamond (BDD, 3 mm diameter). All potentials were measured with respect to the saturated calomel electrode. Rotating disk experiments were conducted using a manually controlled Autolab motor (Ecochemie, Netherlands), connected to a rotating disc electrode. This electrode was a GC (3 mm) electrode and it was rotated over the range 5–25 Hz. 2.3. Synthesis of ferrocene sulfonate compounds 2.3.1. Preparation of ferrocene 1,1 -disulfonic acid A slurry of ferrocene (9.3 g, 50 mmol) in acetic anhydride was stirred at room temperature and 100% sulfuric acid (7.3 g) was added dropwise over a period of 1 h. The internal temperature of the reaction was maintained below 45 ◦ C. The reaction was stirred for a further 2 h at room temperature and the crude product was collected by filtration. The product was washed with acetic anhydride (25 ml) and then pentane (3 × 25 ml) and dried under reduced pressure to give the product as a yellow solid (8.25 g). The disulfonic acid darkened over time and was best stored at low temperature. Due to the apparent instability of the acid product the corresponding salts were prepared. 2.3.2. Preparation of ferrocene 1,1 -disulfonic acid salts The ferrocene 1,1 -disulfonic acid was dissolved in deionised water and neutralized with the corresponding base (such as potassium hydroxide or aqueous ammonia) until the pH of 6–8 was reached. The solution was then freeze dried to give a yellow free flowing powder. The salts were recrystallized from ethanol or water/ethanol mixtures to generate the pure salts. The salts prepared included the potassium, calcium, magnesium, ammonium and tetra-alkyl ammonium variants. 2.3.3. Preparation of ferrocene sulfonic acid A slurry of ferrocene (8.0 g, 43 mmol) in acetic anhydride (60 ml) was stirred at room temperature and 100% chlorosulfonic acid (5.0 g, 43 mmol) was added dropwise over a period of 12 h. The internal temperature of the reaction was maintained below 45 ◦ C. The reaction was stirred for a further 2 h at room temperature and the resulting reaction mixture was poured carefully into iced water (75 ml) and stirred for 1 h. The mixture was then filtered to remove the unreacted ferrocene and the mixture was concentrated in vacuo. The temperature of the solution was maintained at 50 ◦ C or below as higher temperatures decomposed the product. The resulting black solid was extracted with
hot diethyl ether (4 × 50 ml) with the aid of sonication. Removal of the ether under reduced pressure gave the crude product, which was recrystallized from a toluene/pentane mixture to give the product as yellow needle crystals (8.2 g). 3. Results and discussion 3.1. Electrochemical characterization of the ferrocene sulfonate compounds Fig. 1 details the voltammetric response of two solutions containing (A) 1 mM FcSO3 − K+ and (B) 1 mM Fc(SO3 )2 2− 2K+ when dissolved in pH 6.9 phosphate buffer at a GC working electrode at various scan rates (0.01–1 V s−1 ). Analysis of the results in turn reveals an oxidative wave at +0.43 V (versus SCE) with a reduction wave at +0.33 V (versus SCE) for FcSO3 − K+ . The presence of these waves can be attributed to the redox chemistry of the ferrocene moiety. The slightly stretched voltammetric waveshape (peak–peak separation = 100 mV compared to 59.9 mV for fully reversible system) can be attributed to a quasi-reversible electron transfer process (see later). Plots of peak current as a function of the square root of scan rate (not shown) were found to be linear (regression data: √ I (A) = 6.41X + 0.15, where X = scan rate/ (V s−1 ), n = 8, R2 = 0.99), confirming the oxidative and reductive processes are both diffusion controlled. A plot of peak current as a function of concentration of FcSO3 − was found to be linear (regression data: I (A) = 0.006X + 0.335, where X = [FcSO3 − ]/M, n = 5, R2 = 0.997) over the concentration range studied 0–1.1 mM with the peak potentials also being independent of the pH of the solution. Similar voltammetric characteristics were observed for the Fc(SO3 )2 2− 2K+ (Fig. 1B). In this case, a single oxidation wave was observed at +0.65 V (versus SCE) with the corresponding reduction wave at +0.55 V (versus SCE). A comparison of these oxidation potentials with those of FcSO3 − K+ , show that FcSO3 − K+ is easier to oxidize than Fc(SO3 )2 2− 2K+ . The difference in the oxidation potentials of the ferrocene monosulfonate and ferrocene disulfonate ions can be rationalised by the
Fig. 1. Cyclic voltammetric responses of 1 mM Fc(SO3 )− K+ (A) and Fc(SO3 )2 − 2K+ (B) at various scan rates (0.01, 0.025, 0.05, 0.1, 0.2, 0.4, 0.6, 0.8, 1 V s−1 ). Conditions: pH 6.9, phosphate buffer (0.025 M KH2 PO4 , 0.025 M Na2 HPO4 ), GC working electrode.
N.S. Lawrence et al. / Electrochimica Acta 52 (2006) 499–503
Fig. 2. Linear sweep voltammetric responses of (A) 1.8 mM Fc(SO3 )− K+ and (B) 1 mM Fc(SO3 )2 − 2K+ when the working electrode was rotated at various speeds (5, 10, 15, 20 Hz). Conditions: as in Fig. 1, scan rate = 5 mV s−1 .
electron-withdrawing effects of the sulfonate groups; the two sulfonate groups exercise a greater electron-drawing effect and the oxidation of Fe(II) to Fe(III) is therefore more difficult to accomplish [16]. As with FcSO3 − , the slightly stretched voltammetric waveshape (100 mV) and can be attributed to a quasi-reversible electron transfer process occurring (see later). Furthermore, a plot of oxidative and reductive peak currents as a function of the square root of scan rate were found to be linear √ (regression data: I (A) = 1.17X + 0.03, where X = scan rate/ (V s−1 ), n = 9 R2 = 0.995), consistent with a diffusion controlled process. A plot of peak current as a function of concentration of Fc(SO3 )2 2− 2K+ was found to be linear (regression data: I (A) = 0.004X + 0.267, where X = [Fc(SO3 )2 2− ]/M, n = 5, R2 = 0.999) over the concentration range studied 0–1 mM with the peak potentials also being independent of the pH of the solution. To probe the electrochemical system further a rotating disk study was conducted to determine the diffusion coefficient of each species in turn. Fig. 2 details the linear sweep voltammetric responses (scan rate = 5 mV s−1 ) of (A) 1.8 mM FcSO3 − K+ and (B) 1 mM Fc(SO3 )2 2− 2K+ (pH 6.9) at a GC electrode when it was rotated at various speeds (5–20 Hz). Plots of limiting current as a function of the square root of rotation speed were found to be linear in each case. Analysis of these data via the Levich equation for a rotating disk electrode [24] yielded values of 7.1 × 10−6 and 6.9 × 10−6 cm2 s−1 for the diffusion coefficients of the anions FcSO3 − and Fc(SO3 )2 2− , respectively, as potassium salts. Further analysis of the rotating disk data was conducted via Tafel analysis [25] whereby graphs of elec−1 trode potential against log10 (I−1 − Ilim ) were made. In each case, these were found to be linear with gradients of 65 and 71 mV/decade for the FcSO3 − K+ and Fc(SO3 )2 2− 2K+ species, respectively. These values indicate a slight deviation from an ideal electrochemically reversible process occurring in agreement with the cyclic voltammetric data detailed above. To show the generic nature of the electrochemistry of these soluble ferrocene compounds the cyclic voltammetric responses of the ammonium and magnesium salts of the disulfonated species were next examined at a GC working electrode. The cyclic voltammetric response of (A) Fc(SO3 )2 2− (NH4 + )2
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(1 mM) and (B) Fc(SO3 )2 2− (Mg2+ ) (1 mM) at various scan rates (0.01–1 V s−1 ) when dissolved in pH 7 phosphate buffer was next studied. In each case, an increase in the oxidative (+0.66 V (versus SCE)) and reductive (+0.54 V) peak currents was obtained as the scan rate was increased. A plot of peak current as a function of the square root of scan rate was found to be linear for each species consistent with diffusion controlled reactions. A rotating disk study of Fc(SO3 )2 2− (NH4 + )2 (1 mM) over the range of rotation speeds 5–25 Hz was next undertaken. The resulting data (not shown) yielded a linear plot of limiting current as a function of the square root of rotation speed. Analysis of these plots produced value of 6.0 × 10−6 cm2 s−1 for the diffusion coefficient of the anion Fc(SO3 )2 2− as the ammonium salt. Furthermore, a Tafel analysis of the linear sweep voltammogram obtained when the rotation speed was 10 Hz produced a value of 85 mV/decade. This is consistent with the data obtained above for the potassium salt. Finally, the voltammetric signal of Fc(SO3 )2 2− (NH4 + )2 was examined at a BDD electrode, where the electrochemical signal (not shown) produced an oxidative wave at +0.64 V (versus SCE) and a corresponding reduction wave at +0.56 V (versus SCE). These values consistent with those obtained at the GC electrode and suggest that the electron transfer rate is not affected by the underlying electrode substrate. Furthermore, a plot of peak current as a function of the square root of scan rate (0.01–1 V s−1 ) was found to be linear in agreement with the data detailed earlier. Next the voltammetric response of FcSO3 − K+ was studied in the absence and presence of sulfide as a means of examining the possibility of utilizing this species as an electrocatalytic mediator for sulfide detection. The FcSO3 − K+ was chosen in preference to Fc(SO3 )2 2− 2K+ due to its lower oxidation potential. 3.2. Voltammetric response of FcSO3 − to sulfide Fig. 3A details the cyclic voltammetry of a solution containing 1 mM FcSO3 − K+ towards increasing additions of sulfide (0–1 mM, pH 6.9). As can be seen the response of in the absence
Fig. 3. (A) Cyclic voltammetric responses of 1 mM FcSO3 − K+ towards increasing concentrations of sulfide (0–800 M). Conditions: as in Fig. 1. (B) A plot of the increase in oxidation peak current as a function of the sulfide concentration obtained at various pH values.
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of sulfide is consistent with the data detailed in Fig. 1A. Upon the addition of sulfide to the solution an increase in the oxidative peak current is observed at +0.43 V (versus SCE) over the entire concentration range studied, with a corresponding decrease in the reductive current at +0.33 V (versus SCE). This voltammetric response is entirely consistent with the soluble FcSO3 − K+ undergoing an electrocatalytic EC reaction process. The sulfide chemically reduces back the newly formed ferricenium moiety such that it can be regenerated at the electrode surface, thereby producing an enhancement in the oxidative current. These results are analogous to that observed previously for the electrocatalytic reaction of sulfide with ferricyanide [18–22]. As the FcSO3 − K+ acts as a homogeneous mediator for the electrocatalytic oxidation of sodium sulfide, chronoamperometric experiments can be conducted to deduce the catalytic rate constant. The results should follow the chronoamaperometric equation [26,27]: IC λ 0.5 0.5 0.5 π exp(rF (λ ) + exp − 0.5 =λ IL λ where IC is the catalytic current, IL the diffusion limited current and λ = kc0 t (k is the catalytic rate constant and c0 is the bulk initial concentration of sulfide) is the argument of the error function. In the cases, where IC /IL is greater than 1.5, the error function is almost equal to one and the reaction zone is the pure kinetic region. Therefore: IC 0.5 = π0.5 λ0.5 = π0.5 (kc0 t) IL and the rate constant, k, can be deduced. Chronoamperometric experiments for a solution containing 50 M FcSO3 − K+ were conducted both in the absence and presence of sulfide (1–1.4 mM). In these cases, the potential was stepped from 0 to +0.6 V (versus SCE) and the current decay measured. The current was found to decay slowly in the presence of sulfide due to the catalytic reaction process occurring. The currents under each condition was measured at t = 80 ms where IC /IL is greater than 1.5. The catalytic rate constant was found to be equal to 9.52 (±0.81) × 10−3 M s−1 . Next the effect of pH on the analytical response of FcSO3 − K+ towards sulfide was examined. In this case, 1 mM FcSO3 − K+ was dissolved successively in pH 4 (potassium phthalate, 0.05 M), pH 7 (disodium hydrogen phosphate, 0.025 M, potassium dihydrogen phosphate, 0.025 M) and pH 9 (disodium tetraborate, 0.05 M) buffers and the voltammetric response recorded (100 mV s−1 ) in the absence and presence of increasing additions of sulfide (0–800 M). Fig. 3B details the plots of the increase in oxidative peak current as a function of sulfide concentration for each of the pH values studied. It can be seen that as the pH is increased from 4 to 9, the sensitivity increases. This can be rationalized as follows: at pH’s below 6.88 the sulfide is in the form of H2 S [3] and therefore, in order for the ferricenium species to oxidize the H2 S, two protons will have to be removed and the overall reaction rate will be slow compared to the timescale of the voltammetric experiment. At a pH above, the pKa value of H2 S, the sulfide is only singly protonated and as such ‘easier’ to oxidize. Analysis of the voltammetric response
Fig. 4. Cyclic voltammetric responses of 200 M FcSO3 − K+ towards increasing concentrations of (A) sulfide (0–100 M) and (B) mercaptoethanol (0–160 M). Conditions: as in Fig. 1, BDD working electrode.
of sulfide in the absence of the ferrocene sulfonate, at each pH value studied, revealed how varying the pH changed the potential for the direct oxidation of sulfide. Lowering the pH from 9 to 4 increased the oxidation peak potential from 0.6 V (versus SCE) to 1 V (versus SCE). Finally, the voltammetric response of FcSO3 − K+ towards low concentrations of sulfide was examined. Fig. 4 details the voltammetric response of 200 M FcSO3 − towards increasing concentrations of both (A) sulfide (0–100 M) and (B) mercaptoethanol (0–160 M) obtained in pH 6.9 phosphate buffer with a BDD electrode. Analysis of each response in turn first reveals that the use of a BDD electrode has no affect on the voltammetric response of FcSO3 − K+ with an oxidative wave recorded at +0.43 V (versus SCE) and a corresponding reduction wave at +0.33 V (versus SCE) consistent with that obtained at the GC electrode (see above). As shown in Fig. 4A, upon the introduction of sulfide to the solution an increase in the oxidative current is observed with a corresponding decrease in the reduction current. These results are entirely consistent with those detailed in Fig. 3A, and show that the FcSO3 − K+ can be used as a mediator to detect low concentrations of sulfide. A plot of the increase in oxidation peak current as a function of sulfide was found to be linear over the concentration range studied 0–140 M (regression data: I (A) = 0.015X − 0.102, where X = [sulfide]/M, n = 7, R2 = 0.99) with a calculated limit of detection (based on 3sb ) of 14 M. It can be envisaged that the use of pulse voltammetry would lower the detection limit further so it is comparable with those obtained for other electrocatalytic detection methods of sulfide [18–22]. Fig. 4B details the response of FcSO3 − towards increasing additions of mercaptoethanol (0–160 M). It can be seen that upon the introduction of mercaptoethanol to the solutions (40 M) an increase in the oxidative peak current with a corresponding decrease in the reduction peak current was observed. The increase in the oxidation current was found to be linear (regression data: I (A) = 0.005X − 0.042, where X = [sulfide]/M, n = 5, R2 = 0.99) producing a limit of detection of 26 M (based on 3sb ). These results show the generic nature of both the electrocatalytic reaction and detection mechanisms towards thiol determination.
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4. Conclusion The above results describe the voltammetric response of both FcSO3 − K+ and Fc(SO3 )2 2− 2K+ in the absence of sulfide. The results show both species produce voltammograms consistent with quasi-reversible diffusion controlled redox reactions. The FcSO3 − K+ species was found to be easier to oxidize than its Fc(SO3 )2 2− 2K+ counterpart, due to the electron withdrawing affect of the sulfonate group upon the Fe centre. Furthermore, the redox chemistry of FcSO3 − K+ was utilized as a means of promoting the detection of sulfide and thiol species via an electrocatalytic pathway. The response towards sulfide was found be consistent with an EC reaction occurring and produced a linear range over the concentration 0.02–1 mM and a limit of detection of 14 M. These results compare favourably with that of previous studies for the electro-catalytic determination of thiol compounds [18–23]. References [1] P. Patnaik, A Comprehensive Guide to Hazardous Properties of Chemical Substances, second ed., Wiley, New York, 1999. [2] N.S. Lawrence, J. Davis, R.G. Compton, Talanta 52 (2000) 711. [3] N.N. Greenwood, A. Earnshaw, Chemistry of the Elements, Pergamon Press, Oxford, 1986, p. 807. [4] M. Garcia-Calzalda, G. Marban, A.B. Fuertes, Anal. Chim. Acta 380 (1999) 39. [5] V. Stanic, T.H. Etsell, A.C. Pierre, R.J. Mikulka, Electrochim. Acta 43 (1998) 2639. [6] P. Jeroschewski, S. Braun, J. Fresenius Anal. Chem. 354 (1996) 169.
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[7] G. Schiavon, G. Zotti, R. Toniolo, G. Bontempelli, Anal. Chem. 67 (1995) 318. [8] N. Batina, I. Ciglenecki, B. Cosovic, Anal. Chim. Acta 267 (1992) 157. [9] N.S. Lawrence, L. Jiang, T.G.J. Jones, R.G. Compton, Anal. Chem. 75 (2003) 2054. [10] V. Weinmayr, J. Am. Chem. Soc. 77 (1955) 3009. [11] G.R. Knox, P.L. Paulson, J. Chem. Soc. 12 (1958) 692. [12] C. Lee, M.H. Lee, Y.L. Kuang, B.S. Moon, S.B. Rhee, Synth. Met. 55 (1993) 1119. [13] S. Mu, J. Kan, Synth. Met. 132 (2002) 29. [14] S. Mu, D. Shan, Y. Yang, Y. Li, Synth. Met. 135–136 (2003) 199. [15] Y. Yang, S. Mu, Biosens. Bioelectron. 21 (2005) 74. [16] E. Liaudet, F. Battaglini, E.J. Calvo, J. Electroanal. Chem. Interface Electrochem. 293 (1990) 55. [17] M. Tian, S. Dong, Electroanalysis 7 (1995) 1063. [18] P. Jeroschewski, K. Haase, A. Trommer, P. Grundler, Fresenius J. Anal. Chem. 346 (1993) 930. [19] P. Jeroschewski, K. Haase, A. Trommer, P. Grundler, Electroanalysis 6 (1994) 769. [20] P. Jeroschewski, C. Steuckart, M. Kuhl, Anal. Chem. 68 (1996) 4351. [21] N.S. Lawrence, M. Thompson, C. Prado, L. Jiang, T.G.J. Jones, R.G. Compton, Electroanalysis 14 (2002) 499. [22] O. Nekrassova, G.D. Allen, N.S. Lawrence, L. Jiang, T.G.J. Jones, R.G. Compton, Electroanalysis 14 (2002) 1464. [23] J.B. Raoof, R. Ojani, M. Kolbadinezhad, Bull. Chem. Soc. Jpn. 78 (2005) 818. [24] A.J. Bard, L.R. Faulkner, Electrochemical Methods Fundamental and Applications, second ed., Wiley, New York, 2000, p. 339. [25] C.M.A. Brett, A.M.O. Brett, Electrochemistry, Principles, Methods and Applications, Oxford University Press, Oxford, 1993, p. 115. [26] Z. Galus, Fundamentals of Electrochemical Analysis, Ellis Horwood Press, New York, 1994, p. 398. [27] Z.N. Gao, J. Zhang, W.Y. Liu, J. Electroanal. Chem. 580 (2005) 9.
Ferrocene sulfonates as electrocatalysts for sulfide detection Nathan S. Lawrence ∗ , Gary J. Tustin, Michael Faulkner 1 , Timothy G.J. Jones Schlumberger Cambridge Research, High Cross, Madingley Road, Cambridge CB3 0EL, UK Received 21 March 2006; received in revised form 16 May 2006; accepted 16 May 2006 Available online 3 July 2006
Abstract The electrochemical characterization of both the mono- and di-substituted forms of ferrocene sulfonate are given. The results show both species produce voltammograms consistent with quasi-irreversible diffusion controlled redox reactions. The FcSO3 − species was found to be easier to oxidize than its Fc(SO3 )2 2− counterpart, due to the electron withdrawing affect of the sulfonate group on the Fe centre. In the presence of sulfide, the voltammetric response of FcSO3 − is shown to be consistent with the occurrence of an electrocatalytic EC reaction. This analytical response was utilized as a means of determining sulfide and was found to be linear over the concentration 0.02–1 mM with a limit of detection of 14 M. © 2006 Elsevier Ltd. All rights reserved. Keywords: Ferrocene sulfonate; Hydrogen sulfide; Electrocatalytic; Mercaptoethanol; Voltammetry
1. Introduction The determination of sulfide species is important to analytical and environmental chemists [1,2]. This is due to the high toxicity of liberated hydrogen sulfide, as it poses a major problem to those who handle and remove sulfide-contaminated products. The high reactivity of sulfide [3] and hence low accumulative potential of the species minimises the risk of everyday exposure to levels that are likely to cause any serious, adverse health effects. However, there are a number of instances in which localised concentrations of the anion may greatly exceed safety limits and pose significant concerns for those required to work within such environments. Although a wide range of protocols are available for the detection of sulfide which encompass techniques, such as chromatography, spectroscopy and solid state devices [2], electrochemical procedures provide some of the more viable options due to their design simplicity, sensitivity and low cost [4–9]. In the following, we examine the possibility of utilizing the soluble ferrocene derivatives ferrocene monosulfonate (FcSO3 − ) and ferrocene disulfonate (Fc(SO3 )2 2− ) as electrochemical mediators for the detection of sulfide. The synthesis
∗
Corresponding author. Tel.: +44 1223 325224. E-mail address: [email protected] (N.S. Lawrence). 1 Present address: Castle Sixth Form Centre, Kenilworth School, Rouncil Lane, Kenilworth CV8 1FN, UK. 0013-4686/$ – see front matter © 2006 Elsevier Ltd. All rights reserved. doi:10.1016/j.electacta.2006.05.031
of both the mono- and di-substituted ferrocene sulfonates has been described previously [10,11]. Ferrocene sulfonates have been used extensively as anion dopants in the formation of various cationic polymers (for example, polythiophene, polypyrrole [12] and polyaniline [13–15]) or as mediators in enzymatic reactions [16,17]. It can be envisaged that these soluble ferrocene compounds can undergo an electrocatalytic reaction with sulfide and thiol species in an analogous manner to that of the ferrocyanide/ferricyanide redox couple [18–22] and solid ferrocene [23]. In this case, the ferrocene is first oxidized to the ferricenium ion, which is subsequently homogeneously reduced back to ferrocene by the thiol species. This regenerated ferrocence can then be again oxidized at the electrode surface to produce the analytical signal. Herein, we discuss the electrochemistry of various ferrocene sulfonate salts and their subsequent response to hydrogen sulfide. 2. Experimental 2.1. Reagents All chemicals were supplied by Aldrich and used without further purification. The buffered solutions were prepared as follows: pH 4 potassium phthalate buffer (0.05 M), pH 7 disodium hydrogen phosphate (0.025 M) and potassium dihy-
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drogen phosphate (0.025 M) and pH 9 disodium tetraborate (0.05 M). 2.2. Apparatus Electrochemical measurements were recorded using an PGSTAT30 potentiostat (Ecochemie, Netherlands) with a standard three-electrode configuration. A platinum wire (1 mm diameter, Goodfellow Metals, Cambridge, UK) provided the counter electrode and a saturated calomel electrode (SCE, Radiometer, Copenhagen) acted as the reference. The working electrode was composed of either glassy carbon (GC, 2 mm diameter) or boron-doped diamond (BDD, 3 mm diameter). All potentials were measured with respect to the saturated calomel electrode. Rotating disk experiments were conducted using a manually controlled Autolab motor (Ecochemie, Netherlands), connected to a rotating disc electrode. This electrode was a GC (3 mm) electrode and it was rotated over the range 5–25 Hz. 2.3. Synthesis of ferrocene sulfonate compounds 2.3.1. Preparation of ferrocene 1,1 -disulfonic acid A slurry of ferrocene (9.3 g, 50 mmol) in acetic anhydride was stirred at room temperature and 100% sulfuric acid (7.3 g) was added dropwise over a period of 1 h. The internal temperature of the reaction was maintained below 45 ◦ C. The reaction was stirred for a further 2 h at room temperature and the crude product was collected by filtration. The product was washed with acetic anhydride (25 ml) and then pentane (3 × 25 ml) and dried under reduced pressure to give the product as a yellow solid (8.25 g). The disulfonic acid darkened over time and was best stored at low temperature. Due to the apparent instability of the acid product the corresponding salts were prepared. 2.3.2. Preparation of ferrocene 1,1 -disulfonic acid salts The ferrocene 1,1 -disulfonic acid was dissolved in deionised water and neutralized with the corresponding base (such as potassium hydroxide or aqueous ammonia) until the pH of 6–8 was reached. The solution was then freeze dried to give a yellow free flowing powder. The salts were recrystallized from ethanol or water/ethanol mixtures to generate the pure salts. The salts prepared included the potassium, calcium, magnesium, ammonium and tetra-alkyl ammonium variants. 2.3.3. Preparation of ferrocene sulfonic acid A slurry of ferrocene (8.0 g, 43 mmol) in acetic anhydride (60 ml) was stirred at room temperature and 100% chlorosulfonic acid (5.0 g, 43 mmol) was added dropwise over a period of 12 h. The internal temperature of the reaction was maintained below 45 ◦ C. The reaction was stirred for a further 2 h at room temperature and the resulting reaction mixture was poured carefully into iced water (75 ml) and stirred for 1 h. The mixture was then filtered to remove the unreacted ferrocene and the mixture was concentrated in vacuo. The temperature of the solution was maintained at 50 ◦ C or below as higher temperatures decomposed the product. The resulting black solid was extracted with
hot diethyl ether (4 × 50 ml) with the aid of sonication. Removal of the ether under reduced pressure gave the crude product, which was recrystallized from a toluene/pentane mixture to give the product as yellow needle crystals (8.2 g). 3. Results and discussion 3.1. Electrochemical characterization of the ferrocene sulfonate compounds Fig. 1 details the voltammetric response of two solutions containing (A) 1 mM FcSO3 − K+ and (B) 1 mM Fc(SO3 )2 2− 2K+ when dissolved in pH 6.9 phosphate buffer at a GC working electrode at various scan rates (0.01–1 V s−1 ). Analysis of the results in turn reveals an oxidative wave at +0.43 V (versus SCE) with a reduction wave at +0.33 V (versus SCE) for FcSO3 − K+ . The presence of these waves can be attributed to the redox chemistry of the ferrocene moiety. The slightly stretched voltammetric waveshape (peak–peak separation = 100 mV compared to 59.9 mV for fully reversible system) can be attributed to a quasi-reversible electron transfer process (see later). Plots of peak current as a function of the square root of scan rate (not shown) were found to be linear (regression data: √ I (A) = 6.41X + 0.15, where X = scan rate/ (V s−1 ), n = 8, R2 = 0.99), confirming the oxidative and reductive processes are both diffusion controlled. A plot of peak current as a function of concentration of FcSO3 − was found to be linear (regression data: I (A) = 0.006X + 0.335, where X = [FcSO3 − ]/M, n = 5, R2 = 0.997) over the concentration range studied 0–1.1 mM with the peak potentials also being independent of the pH of the solution. Similar voltammetric characteristics were observed for the Fc(SO3 )2 2− 2K+ (Fig. 1B). In this case, a single oxidation wave was observed at +0.65 V (versus SCE) with the corresponding reduction wave at +0.55 V (versus SCE). A comparison of these oxidation potentials with those of FcSO3 − K+ , show that FcSO3 − K+ is easier to oxidize than Fc(SO3 )2 2− 2K+ . The difference in the oxidation potentials of the ferrocene monosulfonate and ferrocene disulfonate ions can be rationalised by the
Fig. 1. Cyclic voltammetric responses of 1 mM Fc(SO3 )− K+ (A) and Fc(SO3 )2 − 2K+ (B) at various scan rates (0.01, 0.025, 0.05, 0.1, 0.2, 0.4, 0.6, 0.8, 1 V s−1 ). Conditions: pH 6.9, phosphate buffer (0.025 M KH2 PO4 , 0.025 M Na2 HPO4 ), GC working electrode.
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Fig. 2. Linear sweep voltammetric responses of (A) 1.8 mM Fc(SO3 )− K+ and (B) 1 mM Fc(SO3 )2 − 2K+ when the working electrode was rotated at various speeds (5, 10, 15, 20 Hz). Conditions: as in Fig. 1, scan rate = 5 mV s−1 .
electron-withdrawing effects of the sulfonate groups; the two sulfonate groups exercise a greater electron-drawing effect and the oxidation of Fe(II) to Fe(III) is therefore more difficult to accomplish [16]. As with FcSO3 − , the slightly stretched voltammetric waveshape (100 mV) and can be attributed to a quasi-reversible electron transfer process occurring (see later). Furthermore, a plot of oxidative and reductive peak currents as a function of the square root of scan rate were found to be linear √ (regression data: I (A) = 1.17X + 0.03, where X = scan rate/ (V s−1 ), n = 9 R2 = 0.995), consistent with a diffusion controlled process. A plot of peak current as a function of concentration of Fc(SO3 )2 2− 2K+ was found to be linear (regression data: I (A) = 0.004X + 0.267, where X = [Fc(SO3 )2 2− ]/M, n = 5, R2 = 0.999) over the concentration range studied 0–1 mM with the peak potentials also being independent of the pH of the solution. To probe the electrochemical system further a rotating disk study was conducted to determine the diffusion coefficient of each species in turn. Fig. 2 details the linear sweep voltammetric responses (scan rate = 5 mV s−1 ) of (A) 1.8 mM FcSO3 − K+ and (B) 1 mM Fc(SO3 )2 2− 2K+ (pH 6.9) at a GC electrode when it was rotated at various speeds (5–20 Hz). Plots of limiting current as a function of the square root of rotation speed were found to be linear in each case. Analysis of these data via the Levich equation for a rotating disk electrode [24] yielded values of 7.1 × 10−6 and 6.9 × 10−6 cm2 s−1 for the diffusion coefficients of the anions FcSO3 − and Fc(SO3 )2 2− , respectively, as potassium salts. Further analysis of the rotating disk data was conducted via Tafel analysis [25] whereby graphs of elec−1 trode potential against log10 (I−1 − Ilim ) were made. In each case, these were found to be linear with gradients of 65 and 71 mV/decade for the FcSO3 − K+ and Fc(SO3 )2 2− 2K+ species, respectively. These values indicate a slight deviation from an ideal electrochemically reversible process occurring in agreement with the cyclic voltammetric data detailed above. To show the generic nature of the electrochemistry of these soluble ferrocene compounds the cyclic voltammetric responses of the ammonium and magnesium salts of the disulfonated species were next examined at a GC working electrode. The cyclic voltammetric response of (A) Fc(SO3 )2 2− (NH4 + )2
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(1 mM) and (B) Fc(SO3 )2 2− (Mg2+ ) (1 mM) at various scan rates (0.01–1 V s−1 ) when dissolved in pH 7 phosphate buffer was next studied. In each case, an increase in the oxidative (+0.66 V (versus SCE)) and reductive (+0.54 V) peak currents was obtained as the scan rate was increased. A plot of peak current as a function of the square root of scan rate was found to be linear for each species consistent with diffusion controlled reactions. A rotating disk study of Fc(SO3 )2 2− (NH4 + )2 (1 mM) over the range of rotation speeds 5–25 Hz was next undertaken. The resulting data (not shown) yielded a linear plot of limiting current as a function of the square root of rotation speed. Analysis of these plots produced value of 6.0 × 10−6 cm2 s−1 for the diffusion coefficient of the anion Fc(SO3 )2 2− as the ammonium salt. Furthermore, a Tafel analysis of the linear sweep voltammogram obtained when the rotation speed was 10 Hz produced a value of 85 mV/decade. This is consistent with the data obtained above for the potassium salt. Finally, the voltammetric signal of Fc(SO3 )2 2− (NH4 + )2 was examined at a BDD electrode, where the electrochemical signal (not shown) produced an oxidative wave at +0.64 V (versus SCE) and a corresponding reduction wave at +0.56 V (versus SCE). These values consistent with those obtained at the GC electrode and suggest that the electron transfer rate is not affected by the underlying electrode substrate. Furthermore, a plot of peak current as a function of the square root of scan rate (0.01–1 V s−1 ) was found to be linear in agreement with the data detailed earlier. Next the voltammetric response of FcSO3 − K+ was studied in the absence and presence of sulfide as a means of examining the possibility of utilizing this species as an electrocatalytic mediator for sulfide detection. The FcSO3 − K+ was chosen in preference to Fc(SO3 )2 2− 2K+ due to its lower oxidation potential. 3.2. Voltammetric response of FcSO3 − to sulfide Fig. 3A details the cyclic voltammetry of a solution containing 1 mM FcSO3 − K+ towards increasing additions of sulfide (0–1 mM, pH 6.9). As can be seen the response of in the absence
Fig. 3. (A) Cyclic voltammetric responses of 1 mM FcSO3 − K+ towards increasing concentrations of sulfide (0–800 M). Conditions: as in Fig. 1. (B) A plot of the increase in oxidation peak current as a function of the sulfide concentration obtained at various pH values.
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of sulfide is consistent with the data detailed in Fig. 1A. Upon the addition of sulfide to the solution an increase in the oxidative peak current is observed at +0.43 V (versus SCE) over the entire concentration range studied, with a corresponding decrease in the reductive current at +0.33 V (versus SCE). This voltammetric response is entirely consistent with the soluble FcSO3 − K+ undergoing an electrocatalytic EC reaction process. The sulfide chemically reduces back the newly formed ferricenium moiety such that it can be regenerated at the electrode surface, thereby producing an enhancement in the oxidative current. These results are analogous to that observed previously for the electrocatalytic reaction of sulfide with ferricyanide [18–22]. As the FcSO3 − K+ acts as a homogeneous mediator for the electrocatalytic oxidation of sodium sulfide, chronoamperometric experiments can be conducted to deduce the catalytic rate constant. The results should follow the chronoamaperometric equation [26,27]: IC λ 0.5 0.5 0.5 π exp(rF (λ ) + exp − 0.5 =λ IL λ where IC is the catalytic current, IL the diffusion limited current and λ = kc0 t (k is the catalytic rate constant and c0 is the bulk initial concentration of sulfide) is the argument of the error function. In the cases, where IC /IL is greater than 1.5, the error function is almost equal to one and the reaction zone is the pure kinetic region. Therefore: IC 0.5 = π0.5 λ0.5 = π0.5 (kc0 t) IL and the rate constant, k, can be deduced. Chronoamperometric experiments for a solution containing 50 M FcSO3 − K+ were conducted both in the absence and presence of sulfide (1–1.4 mM). In these cases, the potential was stepped from 0 to +0.6 V (versus SCE) and the current decay measured. The current was found to decay slowly in the presence of sulfide due to the catalytic reaction process occurring. The currents under each condition was measured at t = 80 ms where IC /IL is greater than 1.5. The catalytic rate constant was found to be equal to 9.52 (±0.81) × 10−3 M s−1 . Next the effect of pH on the analytical response of FcSO3 − K+ towards sulfide was examined. In this case, 1 mM FcSO3 − K+ was dissolved successively in pH 4 (potassium phthalate, 0.05 M), pH 7 (disodium hydrogen phosphate, 0.025 M, potassium dihydrogen phosphate, 0.025 M) and pH 9 (disodium tetraborate, 0.05 M) buffers and the voltammetric response recorded (100 mV s−1 ) in the absence and presence of increasing additions of sulfide (0–800 M). Fig. 3B details the plots of the increase in oxidative peak current as a function of sulfide concentration for each of the pH values studied. It can be seen that as the pH is increased from 4 to 9, the sensitivity increases. This can be rationalized as follows: at pH’s below 6.88 the sulfide is in the form of H2 S [3] and therefore, in order for the ferricenium species to oxidize the H2 S, two protons will have to be removed and the overall reaction rate will be slow compared to the timescale of the voltammetric experiment. At a pH above, the pKa value of H2 S, the sulfide is only singly protonated and as such ‘easier’ to oxidize. Analysis of the voltammetric response
Fig. 4. Cyclic voltammetric responses of 200 M FcSO3 − K+ towards increasing concentrations of (A) sulfide (0–100 M) and (B) mercaptoethanol (0–160 M). Conditions: as in Fig. 1, BDD working electrode.
of sulfide in the absence of the ferrocene sulfonate, at each pH value studied, revealed how varying the pH changed the potential for the direct oxidation of sulfide. Lowering the pH from 9 to 4 increased the oxidation peak potential from 0.6 V (versus SCE) to 1 V (versus SCE). Finally, the voltammetric response of FcSO3 − K+ towards low concentrations of sulfide was examined. Fig. 4 details the voltammetric response of 200 M FcSO3 − towards increasing concentrations of both (A) sulfide (0–100 M) and (B) mercaptoethanol (0–160 M) obtained in pH 6.9 phosphate buffer with a BDD electrode. Analysis of each response in turn first reveals that the use of a BDD electrode has no affect on the voltammetric response of FcSO3 − K+ with an oxidative wave recorded at +0.43 V (versus SCE) and a corresponding reduction wave at +0.33 V (versus SCE) consistent with that obtained at the GC electrode (see above). As shown in Fig. 4A, upon the introduction of sulfide to the solution an increase in the oxidative current is observed with a corresponding decrease in the reduction current. These results are entirely consistent with those detailed in Fig. 3A, and show that the FcSO3 − K+ can be used as a mediator to detect low concentrations of sulfide. A plot of the increase in oxidation peak current as a function of sulfide was found to be linear over the concentration range studied 0–140 M (regression data: I (A) = 0.015X − 0.102, where X = [sulfide]/M, n = 7, R2 = 0.99) with a calculated limit of detection (based on 3sb ) of 14 M. It can be envisaged that the use of pulse voltammetry would lower the detection limit further so it is comparable with those obtained for other electrocatalytic detection methods of sulfide [18–22]. Fig. 4B details the response of FcSO3 − towards increasing additions of mercaptoethanol (0–160 M). It can be seen that upon the introduction of mercaptoethanol to the solutions (40 M) an increase in the oxidative peak current with a corresponding decrease in the reduction peak current was observed. The increase in the oxidation current was found to be linear (regression data: I (A) = 0.005X − 0.042, where X = [sulfide]/M, n = 5, R2 = 0.99) producing a limit of detection of 26 M (based on 3sb ). These results show the generic nature of both the electrocatalytic reaction and detection mechanisms towards thiol determination.
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4. Conclusion The above results describe the voltammetric response of both FcSO3 − K+ and Fc(SO3 )2 2− 2K+ in the absence of sulfide. The results show both species produce voltammograms consistent with quasi-reversible diffusion controlled redox reactions. The FcSO3 − K+ species was found to be easier to oxidize than its Fc(SO3 )2 2− 2K+ counterpart, due to the electron withdrawing affect of the sulfonate group upon the Fe centre. Furthermore, the redox chemistry of FcSO3 − K+ was utilized as a means of promoting the detection of sulfide and thiol species via an electrocatalytic pathway. The response towards sulfide was found be consistent with an EC reaction occurring and produced a linear range over the concentration 0.02–1 mM and a limit of detection of 14 M. These results compare favourably with that of previous studies for the electro-catalytic determination of thiol compounds [18–23]. References [1] P. Patnaik, A Comprehensive Guide to Hazardous Properties of Chemical Substances, second ed., Wiley, New York, 1999. [2] N.S. Lawrence, J. Davis, R.G. Compton, Talanta 52 (2000) 711. [3] N.N. Greenwood, A. Earnshaw, Chemistry of the Elements, Pergamon Press, Oxford, 1986, p. 807. [4] M. Garcia-Calzalda, G. Marban, A.B. Fuertes, Anal. Chim. Acta 380 (1999) 39. [5] V. Stanic, T.H. Etsell, A.C. Pierre, R.J. Mikulka, Electrochim. Acta 43 (1998) 2639. [6] P. Jeroschewski, S. Braun, J. Fresenius Anal. Chem. 354 (1996) 169.
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