Interaction between Silver and Other Heavy Metal Ions and the

School of Chemistry, La Trobe University, Melbourne, Victoria 3083, Australia ... Publication Date (Web): August 15, 1997. Copyright © 1997 American ...
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Anal. Chem. 1997, 69, 3353-3359

Interaction between Silver and Other Heavy Metal Ions and the Ionophore S,S-Dipropyl Pyridine-2,6-dicarbothioate Robert W. Cattrall,* C. Grace Gregorio, and Richard D. Webster

School of Chemistry, La Trobe University, Melbourne, Victoria 3083, Australia Alan M. Bond

Department of Chemistry, Monash University, Melbourne, Victoria 3168, Australia Keith B. Oldham

Department of Chemistry, Trent University, Peterborough, Ontario K9J 7B8, Canada

A detailed study has been made of the interaction of the neutral metal ion carrier S,S-dipropyl pyridine-2,6-dicarbothioate in methanol solution with silver perchlorate. Evidence has been found by potentiometry and confirmed by electrospray mass spectrometry for the existence of both 1:1 and 1:2 metal/ligand complexes. The formation constants of the complexes in methanol have been determined from the potentiometry data and have been shown to be low (K1 ≈ 50 L mol-1, K2 ≈ 8 L mol-1). The value for K1 has been confirmed by NMR and voltammetric measurements. The high Ag+ selectivity found for the ionophore in a PVC-based sensor implies that the formation constants for the ionophore with other metal ions must be even lower than that for the ionophore with silver. A chemical and voltammetric study of the interaction between the ionophore and certain metal ions in methanol has revealed that the ionophore structure breaks down in the presence of Ag+, Hg2+, and Pb2+. The Hg2+ ion is the most reactive of the three ions and rapidly leads to cleavage of the C(O)-S bonds in the ionophore, with the subsequent precipitation of a product containing the thiol fragment. This reaction explains the irreversible poisoning effect that the Hg2+ ion has on the silver response of the sensor. In contrast, the decomposition reaction in the presence of the Ag+ ion is extremely slow, which explains why the short-term silver response of the sensor is excellent. However, the reaction has severe implications with respect to the lifetime of the sensor. A mechanism is proposed for the ionophore decomposition reaction in the presence of metal ions. We have recently reported1 a highly selective polymer membrane potentiometric sensor for silver which is based on a new ionophore, S,S-dipropyl pyridine-2,6-dicarbothioate (Figure 1). The use of sulfur-containing ionophores in polymer membrane potentiometric chemical sensors for silver ions has been reported (1) Bates, M. R. M.; Cardwell, T. J.; Cattrall, R. W.; Deady, L. W.; Gregorio, C. G. Talanta 1995, 42, 999. S0003-2700(97)00008-5 CCC: $14.00

© 1997 American Chemical Society

Figure 1. S,S-Dipropyl pyridine-2,6-dicarbothioate.

frequently,2,3 and their affinity for ions like silver is associated with the donor power of the soft basic sulfur atom for the soft acidic metal ion. Typically, the thioic S-esters display much higher complexing affinities for silver than for other class b acceptor ions,1 such as Pb2+ and Cd2+, which is seen in the high selectivity toward silver of sensor membranes made with these reagents. However, the sensors suffer from severe interference from the Hg2+ ion, which irreversibly destroys their response.1,3 This study was undertaken to investigate the interaction between the thioic S-ester and Ag+ as well as other heavy metals, in particular Hg2+ because of its strong poisoning effect. It was also of interest to obtain information regarding the formation constant of the complex with silver in an attempt to relate it to the high selectivity of the sensor for this ion. Formation constant studies have been carried out on a number of ionophores used in potentiometric sensors; in particular, the naturally occurring macrotetrolide antibiotics such as valinomycin,4 nigecirin,5 monensin,5 and nonactin6 have all been studied (2) (a) Lai, M. T.; Shih, J. S. Analyst 1986, 111, 891. (b) Oue, M.; Kimura, K.; Akama, K.; Tanaka, M.; Shono, T. Chem. Lett. 1988, 409. (c) Oue, M.; Akama, K.; Kimura, K.; Tanaka, M.; Shono, T. J. Chem. Soc., Perkin Trans. 1 1989, 1675. (d) O’Connor, K. M.; Svehla, G.; Harris, S. J.; Mc-Kervey, M. A. Anal. Proc. 1992, 30, 137. (e) O’Connor, K. M.; Svehla, G.; Harris, S. J.; Mc-Kervey, M. A. Talanta 1993, 39, 1549. (f) Brzozka, Z.; Cobben, P. L. H.; Reinhouldt, D.; Edema, J. J. H.; Butler, J.; Kellogg, R. M. Anal. Chim. Acta 1993, 273, 139. (3) (a) Casabo´, J.; Flor, T.; Romero, M. I.; Teixidor, F.; Pe´rez-Jimenez, C. Anal. Chim. Acta 1994, 294, 207. (b) Casabo´, J.; Teixidor, F.; Escriche, L.; Vin ˜as, C.; Pe´rez-Jimenez, C. Adv. Mater. 1995, 7, 238. (4) (a) Fru ¨ h, P. U.; Clerc, J. T.; Simon, W. Helv. Chim. Acta 1971, 54, 1445. (b) Wipf, H. K.; Pioda, A. R.; Stefanac, Z; Simon, W. Helv. Chim. Acta 1968, 51, 377. (5) Lutz, W. K.; Wipf, H. K.; Simon, W. Helv. Chim. Acta 1970, 53, 1741.

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extensively, and the results have been useful in understanding the selective binding of the alkali metal ions by these molecules. Similar studies on complexes with synthetic molecules used in sensors, such as the the alkali and alkaline earth cryptates7 and complexes with cyclic polyethers,8 have allowed selectivities to be deduced from the formation constants. Methods used for the determination of the formation constants of complexes of the types mentioned invariably employ solvents such as methanol or ethanol because of solubility constraints in aqueous media. The values obtained for the formation constants are thus only a guide to the complexing power of an ionophore for a particular ion in a sensor membrane, which is a very different environment from that of the solvent. Nevertheless, it is the relative complexing power of the ionophore for different ions which is important in determining selectivity, and qualitative information on this can be obtained from the determination of the formation constants. Many techniques are available for the determination of formation constants in solution, but electrochemical methods such as potentiometry predominate. One interesting development is the work of Bakker et al.,9 who have introduced a novel method for the determination of the formation constant of a metal ion/ ionophore complex in an optode membrane. Such a medium, of course, mimics the environment in a potentiometric sensor membrane extremely well since the two membrane compositions are almost identical. EXPERIMENTAL SECTION Reagents. S,S-Dipropyl pyridine-2,6-dicarbothioate was synthesized as described previously.1 The metal perchlorate salts were all analytical grade and were obtained from Aldrich (Milwaukee, WI). The tetraalkylammonium salts, But4NClO4, Et4NBF4 and Et4NPF6 were of electrochemical grade from Southwestern Analytical (Austin, TX). Et4NClO4 was obtained from GFS Chemicals (Columbus, OH). Potassium tetrakis(4-chlorophenyl)borate was 98% pure from Aldrich. HPLC grade methanol was supplied by Mallinckrodt (Paris, KY) and was dried with molecular sieves. Deuterated methanol was obtained from Novachem (Melbourne, Australia). Tetrahydrofuran (THF) from BDH (Poole, UK) was purified by passing through an activated alumina column.1 All aqueous solutions were made using >15 MΩ‚cm NANO-pure water (Barnstead, Dubuque, IA). Electrochemical Measurements. Potentiometric measurements were performed at 25.0 ( 0.05 °C using a thermostated glass cell equipped with inlets for the electrodes, for a N2 bubbler, and for the addition of solid ionophore. The cell was stirred magnetically. An Ag/AgCl electrode was used as the indicator electrode and was conditioned overnight in 0.01 M AgClO4 prior to use. An Orion (Boston, MA) double-junction reference electrode was used with saturated NaCl (saturated with Ag+) as the internal filling solution and 0.1 M Et4NClO4 in methanol as the (6) (a) Zu ¨ st, Ch. U.; Fru ¨ h, P. U.; Simon, W. Helv. Chim. Acta 1973, 56, 495. (b) Pioda, A. R.; Wachter, H. A.; Dohner, R. E.; Simon, W. Helv. Chim. Acta 1967,50, 1373. (7) (a) Lehn, J. M.; Sauvage, J. P. Chem. Commun. 1971, 440. (b) Cheney, J.; Lehn, J. M.; Sauvage, J. P.; Stubbe, M. E. Chem. Commun. 1972, 1100. (8) (a) Frensdorff, H. K. J. Am. Chem. Soc. 1971, 93, 600. (b) Izatt, R. M.; Nelson, D. P.; Rytting, J. H.; Haymore, B. L.; Christensen, J. J. J. Am. Chem. Soc. 1971, 93, 1619. (c) Rechnitz, G. A.; Eyal, E. Anal. Chem. 1972, 44, 370. (d) Michaux, G.; Reisse, J. J. Am. Chem. Soc. 1982, 104, 6895. (9) Bakker, E.; Willer, M.; Lerchi, M.; Seller, K.; Pretsch, E. Anal. Chem. 1994, 66, 516.

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outer filling solution. Potential measurements were made with an Orion Model 901 Ionalyser. The criterion of stability of the potential was that the measurement was considered to have reached a steady value when it did not change by more than ( 0.1 mV over a 2-min period. All solutions used in the study contained 0.1 M Et4NClO4 to maintain a constant ionic strength. Polarograms were recorded with a Metrohm 646 VA processor and a 647 stand with a mercury drop time of 0.8 s. An Ag/Ag+ (0.05 M AgNO3) reference electrode was used which was separated from the test solution by a liquid junction consisting of the solvent/electrolyte solution (e.g., methanol containing 0.1 M tetrabutylammonium perchlorate (Bu4NClO4)) in order to avoid contamination from Ag+ ions. Solutions were thermostated at 25.0 ( 0.2 °C and were degassed with methanol-saturated nitrogen, and a continuous stream of nitrogen was passed over the solution while measurements were made. Linear scan voltammetric measurements were made at 25.0 ( 0.05 °C with a Cypress Systems (Lawrence, KA) Model CS1090 instrument using a Metrohm rotating platinum disk electrode (3 mm diameter) rotated at ∼1000 rpm via the Metrohm rotating disk assembly, Model 628-10. The electrode was frequently polished with a diamond paste (1 µm)/water slurry and rinsed in methanol. The same reference electrode was used as for the potentiometry, and a platinum wire was used as the auxiliary electrode. Solutions were degassed with methanol-saturated nitrogen as for the polarography. NMR Experiments. 1H and 13C NMR spectra were recorded at 298 K on a Bruker AM-300 spectrometer in the appropriate deuterated solvent, CDCl3, (CD)3SO, or CD3OD. Electrospray Mass Spectrometry (ESMS). Mass spectra were obtained with a VG Bio-Q triple-quadrupole mass spectrometer (VG Bio-Tech, Altrincham, UK) using a water/methanol/ acetic acid (50:50:1) mobile phase. Sample solutions in methanol were injected via a Rheodyne injector fitted with a 10-µL loop. A Phoenix 20 LC microsyringe pump delivered the solution to the vaporization nozzle of the instrument at a flow rate of 3 µL min-1. The carrier solution contained 6 µM gramicidin S for tuning the spectrometer. Nitrogen was used both as the drying gas and for nebulization, with flow rates of 3 and 100 mL min-1, respectively. A voltage of 35 V was applied to the first skimmer electrode (B1). Other Instrumental Measurements. Infrared spectra were recorded with a Model 1720X Perkin-Elmer FT-IR spectrophotometer. Using electron ionization (EI) at 70 eV and chemical ionization (CI) with methane as the reagent gas, mass spectra were recorded with a Hewlett-Packard 5988A mass spectrometer with an RTE-A data system. Analyses by atomic absorption spectrometry were carried out by using either a Varian (Melbourne, Australia) Model 275 spectrometer or a GBC Scientific Equipment (Melbourne, Australia) Model 933 spectrometer. Other Analyses. TLC analysis was performed on Merck SiO2 60 F254 precoated aluminium sheets, visualized with UV light. Elemental analyses were performed by the Chemistry Department, University of Otago, Dunedin, New Zealand. RESULTS AND DISCUSSION Previous work1 has shown that the ionophore S,S-dipropyl pyridine-2,6-dicarbothioate (Figure 1), immobilized in a PVC membrane, produces a highly selective sensor for silver ions. Thus, there was an expectation that the compound in the membrane would form a cationic complex with Ag+ with a reasonably high formation constant and that the complex would

Table 1. Formation Constants Obtained by Potentiometry [AgT] (mol L-1)

K1 (L mol-1)

K2 (L mol-1)

6.62 × 10-3 9.93 × 10-3 1.13 × 10-2

52.6 ( 0.2a 49.7 ( 0.1 43.1 ( 1.3

6.8 ( 0.1 8.0 ( 0.5 9.2 ( 0.8

a Standard deviation from curve-fitting (sum of squares about regression).

equation may be rearranged to a quadratic in [LT], the positive solution of which is

[LT] ) 2[AgT] - 2[Ag+] + Figure 2. Total ionophore concentration versus free [Ag+] for [AgT] of (O) 6.62, (]) 9.93, and (0) 11.3 mM. Dashed line, computer fit.

predominately have a 1:1 stoichiometry and be able to exchange readily with silver ions from an aqueous solution. Potentiometry. In order to obtain further insight into the complexation reaction between the ionophore and silver, experiments were carried out to determine the stoichiometry of the complex and to evaluate its formation constant using potentiometry. The potentiometry experiments were conducted in methanol containing 0.1 M Et4NClO4 for three starting silver concentrations, 6.62, 9.93, and 11.3 mM. Each of the silver solutions was spiked by successive weighed increments of the ionophore until its total concentration, [LT], was in 10-fold excess over the total silver concentration, [AgT]. The solution was thoroughly stirred after each addition and the steady state potential recorded. These potential values were used to determine the concentrations, [Ag+], of free silver ion from appropriate calibration graphs. Finally, for each of the three values of the total silver concentration, the total ionophore concentration, [LT], was plotted as a function of the free silver ion concentration, [Ag+], to yield the data points plotted in Figure 2. The experimental data were tested against a model which assumes that both 1:1 and 1:2 metal/ionophore complexes are formed. With L denoting the ionophore, the stepwise formation constants for these complexes are K1 ) [AgL+]/[Ag+][L] and K2 ) [AgL2+]/[AgL+][L]. The total silver and ionophore concentrations in solution are given by [AgT] ) [Ag+] + [AgL+] + [AgL2+] and [LT] ) [L] + [AgL+] + 2[AgL2+], respectively. Combination and rearrangement of these four relationships gives

[LT] - [AgT] + [Ag+] ) [LT] - 2[AgT] + 2[Ag+]

[

]

K1K2[Ag+]([LT] - 2[AgT] + 2[Ag+]) 1 1 + 1 - K1[Ag+] 1 + K1[Ag+]

(1)

For any one sequence of measurements, this is an equation relating the potentiometrically determined free silver ion concentration, [Ag+], to the known total ionophore concentration, [LT], and the known constant total silver concentration, [AgT]. The

[x (

1 - K1[Ag+] 2K2

1+

) ]

4K2 [AgT] - 1 - 1 (2) K1 [Ag+]

Hence, with the correct values of the formation constants K1 and K2, this equation should describe the data in Figure 2, if our model of the chemistry is correct. Iterative nonlinear computer curvefitting was employed to determine the best values of the formation constants to fit the data. The experimental/theory comparison is presented in Figure 2 in the form of graphs of the total ionophore concentration versus the concentration of free Ag+ for the three total silver concentrations. The dashed line in each case is the computer fitted curve, and the corresponding values for K1 and K2 are given in Table 1. The standard deviations from the curve-fitting are also shown in Table 1. A reasonable agreement is obtained for the formation constants for the three total silver concentrations, and the fit to the model as shown in Figure 2 is extremely good. The slight deviation from the model for the highest [AgT] concentration is probably due to the ligand decomposition reaction described later in the paper. The magnitude of the value for K1 is very surprising and unexpected, since such a low formation constant would seem to be contrary to the observation that the ionophore produces a sensor which shows high selectivity for Ag+. Most often, measured formation constants10 for cation-ionophore complexes are of the order of 102-104 M-1. The low value for K2 is less surprising, since it was anticipated that the formation of the 1:1 complex would be favored. Nevertheless, it appears that the ionophore satisfies the requirement of an ion carrier for a potentiometric sensor of having the capability of complexing Ag+ and allowing reversible and rapid ion exchange to occur. The high selectivity of the sensor implies even lower values for the formation constants for complexation between the ionophore and other cations (except for Hg2+, which irreversibly poisons the sensor). Electrospray Mass Spectrometry. Verification of the existence of the two complexes was obtained by electrospray mass spectrometry (ESMS) using solutions of AgClO4 and the ionophore S,S-dipropyl pyridine-2,6-dicarbothioate in methanol. A solution containing a 1:1 mole ratio of ionophore/silver gave peaks in the mass spectrum at m/z 390 and 392, which correspond to (10) Dobler, M. Ionophores and Their Structures; John Wiley and Sons: New York, 1981.

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the 107Ag and 109Ag isotopomers of AgL+. Another set of peaks became apparent at m/z 673 and 675 in addition to the two at m/z 390 and 392 when the ratio of ionophore/metal was increased to 2:1. These additional peaks correspond to the107Ag and 109Ag isotopomers of AgL2+. No higher complexes were observed in the mass spectrum, even when the ionophore/silver ratio was increased to 10:1. NMR. An independent estimate of the magnitude of K1 was obtained from 1H NMR spectroscopy. This was done by measuring the ionophore proton chemical shifts associated with complexation to silver as increasing amounts of the metal ion were added to a known amount of the ionophore. The chemical shifts were measured with respect to the center of the signal for the protons in the pyridine ring, which was in turn referenced to TMS. Three concentrations of the ionophore in deuterated methanol were studied: 0.028, 0.063, and 0.115 M. Initially, the proton spectra were recorded for the three concentrations, and then a weighed amount of AgClO4 was added to each tube, the tube was shaken thoroughly, and the spectrum was recorded. This was repeated until the metal ion was ∼10-fold in excess of the ionophore. These conditions, particularly the final one, precluded the formation of the 2:1 ionophore/silver complex. From the 1H NMR results, it was apparent that the ionophore and its Ag+ complex were undergoing rapid exchange on the NMR time scale. The chemical shift observed, δobs, for the pyridine protons on the addition of silver to the solution is the sum of their chemical shifts in the free and the complexed forms, δo and δc, respectively, weighted according to the relative concentrations of the two forms as given by eq 3.

δobs ) δo([L]/[LT]) + δc([AgL ]/[LT]) +

(3)

Equation 4 can be derived by the inclusion of the expressions for the total ionophore and silver concentrations into the expression for 1/K1.

[AgL+]/K1 ) [AgT][LT] - [LT][AgL+] - [AgT][AgL+] + [AgL+]2 (4) Equation 5 gives the solution to this quadratic equation.

[AgL+] ) ([AgT] + [LT] + 1/K1)/2 ( 1

/2 x([AgT] + [LT] + 1/K1)2 - 4[AgT][LT] (5)

Equation 3 can now be written in the following form (only the negative square-root term of eq 5 is meaningful), where ∆δobs ) (δobs - δo) and ∆δc ) (δc - δo).

∆δobs )

∆δc 2[LT]

[[AgT] + [LT] + 1/K1 -

x([AgT] + [LT] + 1/K1)2 - 4[AgT][LT]]

(6)

Equation 6 contains only two unknown parameters, ∆δc and K1, which are evaluated by iterative computer curve-fitting. Plots of ∆δobs as a function of [AgT] for the three ionophore concentrations in deuterated methanol are shown in Figure 3, together with the best-fit curve (dotted line). The values obtained 3356 Analytical Chemistry, Vol. 69, No. 16, August 15, 1997

Figure 3. ∆δobs versus [AgT] for [LT] of (O) 28.0, (]) 63.0, and (0) 115 mM. Dashed line, computer fit. Table 2. K1 Values from 1H NMR [LT] (mol L-1)

∆δ∞

K1 (L mol-1)

2.8 × 10-2 6.3 × 10-2 1.15 × 10-1

0.209 ( 0.003a 0.208 ( 0.004 0.222 ( 0.007

55 ( 5 71 ( 14 45 ( 13

a Standard deviation from curve-fitting (sum of squares about regression).

for ∆δc and K1 are given in Table 2, which also gives the standard deviations obtained from the computer fit. Even though there are relatively large standard deviations associated with the K1 values, there is agreement with the values obtained from potentiometry, and this serves to support the conclusion that the formation constant for the 1:1 complex is low. Linear Scan Voltammetry. An attempt was made to evaluate the formation constants of the ionophore/silver complexes using voltammetry. For this, it was necessary to use a rotating platinum electrode because of the proximity of the Ag+ reduction potential to that of the dissolution of mercury. The potential window for the reduction was set by initially recording a linear sweep voltammogram of the electrolyte solution (0.1 M Et4NClO4 in methanol). Then, the voltammogram of a AgClO4 solution in the electrolyte was recorded. The reaction studied involves the deposition of metallic silver on a bare platinum surface (Ag+aq + e ) Ags), which is assumed to be reversible. Weighed increments of the ionophore were then added as for the potentiometry, and the voltammogram was recorded after each addition. Three initial silver concentrations, [AgT] ) 0.25, 0.5, and 1.0 mM, were used in the voltammetric work. As in polarography, the half-wave potential, E1/2, of a reversible reduction wave at a rotating disk electrode is predicted to shift with the total ligand concentration, [LT], as described by Cohen et al.11 Unfortunately, although there was a shift in E1/2 of the Ag+ reduction wave as a function of [LT] (seen in Figure 4 for 1 mM [AgT]), as expected when weak complexation occurs with a K1 value of about 40, the large decrease in the limiting current precludes quantitative evaluation of the equilibrium constants. Additionally, the assumption of complete reversibility at all ligand (11) Cohen, S. H.; Reynold, I. T.; Kleinberg, J. J. Am. Chem. Soc. 1960, 82, 1844.

Table 3. Analysis of the Products of the Reaction of the Ionophore with Hg, Pb, and Ag Perchlorates analytical data (% m/m) compound 2, C3H7SClO4Hg 2a, C15H39S5Cl3O12Hg3 3, C18H18N2O16Cl2Pb 4, C15H35S5Cl2O8Ag7

b

calcd found calcd found calcd found calcd found

C

H

N

S

Cl

metal

10.1 9.44 14.1 14.3 27.3 27.1 13.5 13.6

1.78 1.88 2.92 2.51 2.07 2.27 2.63 2.55

nil nil nil nil 3.61 3.52 nil nil

8.46 8.54 12.5 12.6 nil nil 12.1 12.0

9.10 9.44 8.30 7.5-8.5c 8.95 8.90 5.33 5.29

53.5 ndb 47.0 nd 26.0 27.6 56.8 56.5

yielda (mg) 105 164 13.0 nd

a Yield calculated from mass of product obtained from reaction between 200 mg of ionophore and an equimolar amount of metal perchlorate. nd, not determined. c Inteferences present, estimated range.

Figure 4. Voltammograms for [AgT] ) 1.00 mM and [LT] ) (a) 0, (b) 5.0, (c) 10.0, (d) 20, (e) 30, (f) 40, (g) 50, (h) 70, (i) 80, and (j) 100 mM.

concentrations is not met under voltammetric conditions. The decrease in the limiting current is too great to be explained by diffusion coefficient differences of uncomplexed and complexed silver. The significant decrease is most likely associated with the decomposition reactions described below, which are accelerated by the conditions of the voltammetric experiment. Chemical Studies. A chemical study was undertaken to investigate further the interaction between the ionophore and Ag+ as well as the interaction with other heavy metals, in particular Hg2+, because of its strong poisoning effect on the sensor response.1 The study was carried out on both the short (seconds) and long (days) time scales. The short time scale experiments for Hg2+, Pb2+, and Cd2+ were monitored polarographically by measuring the changes in the polarogram for the metal ion reduction processes when an increasing concentration of ionophore was added to a methanol solution containing the metal salt and supporting electrolyte. The long time scale experiments were conducted by mixing equimolar quantities of the ionophore (∼200 mg) with the appropriate metal ion in methanol (20 mL) and leaving the solutions to stand (in the dark) for a period of time at room temperature. After several hours, or days depending on the metal, the products of the reaction, if any, were isolated and identified by NMR spectroscopy, infrared spectroscopy, TLC, and microanalysis. Reaction with Hg2+. When equimolar amounts of the ionophore and Hg(ClO4)2 are mixed together in methanol, a fine white

precipitate forms after several hours which can be redissolved with gentle heating. 1H and 13C NMR experiments indicated that the precipitate contained the S-Pr chain (C3H7S) and no aromatic groups. Microanalysis of the solid indicated the presence of carbon, hydrogen, chlorine, and sulfur but no nitrogen, and atomic absorption spectrometry showed the presence of mercury. The analytical results are given in Table 3 (2). These data, along with the NMR studies, indicate an empirical formula of C3H7SClO4Hg. Although this compound has a simple empirical formula, compounds of this type have been shown (by vibrational spectroscopy) to exist in monomeric, dimeric, and complex polymeric forms in the solid state.12 Further analysis of the mother liquor indicated that two other compounds (in addition to the starting material) were formed in the reaction. These were separated and purified from the reaction mixture by preparative scale column chromatography, identified by NMR spectroscopy, mass spectrometry, infrared spectroscopy and TLC, and shown to be the mono- and dimethyl esters formed by replacement of the thiol groups in the ionophore by methoxy groups. By considering all the products of the reaction, a mechanism is proposed in Figure 5, which involves mercury reacting with one of the sulfur atoms on the ionophore (1) to form 2, with the OCH3 group from methanol simultaneously replacing the thiol groups in the ester to form 1b and 1c. It is possible that the ionophore alone could undergo a hydrolysis-type reaction in methanol solution, with the thioester groups being replaced with methoxy groups, and this was tested. However, it was found that, if the reaction does occur, it is extremely slow at room temperature in an unstirred solution, since 1b and 1c were not detected by TLC 4 weeks after 200 mg of the ionophore was dissolved in 20 mL of methanol. Thus, the conclusion is that mercury acts to catalyze this reaction, presumably by the interaction of Hg with the sulfur atoms weakening the C(O)-S bond, as indicated in Figure 5. The reaction of the ionophore with Hg2+ also was studied by polarography. The addition of 2 molar equiv of the ionophore to a 5 mM solution of Hg(ClO4)2 in methanol containing 0.1 M tetrabutylammonium perchlorate (Bu4NClO4) results immediately in the formation of two new polarographic reduction waves at negative potentials (see Figure 6a). The ionophore is, itself, able (12) (a) Biscarini, P.; Fusina, L.; Nivellini, G. Spectrochim. Acta A 1980, 36, 593. (b) Canty, A. J. Spectrochim. Acta A 1981, 37, 283.

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immediate on the short voltammetric time scale, and an excess of ionophore is needed to bring about complete reaction. These results provide some insight into why the polymer membrane chemical sensor suffers from severe interference from Hg2+, since the ionophore breaks down quite rapidly and irreversibly in the presence of Hg2+ and the sensor ceases to function.1 The mechanism for this is most likely associated with the rapid hydrolysis of the thioester group in the presence of water and Hg2+ to form the carboxylic acid and 2. Polarographic studies of other organic mercury compounds of similar structure14 also commonly display two or three polarographic waves at potentials similar to those shown in Figure 6. At glassy carbon and platinum electrodes, the voltammetric waves corresponding to processes 1 and 2 are also evident, but the voltammetry shows signs of the reduced compound (or other products of the reduction) strongly adsorbing onto the electrode. This is consistent with previous studies15 on mercury dithiolates, where reduction steps associated with the following reactions are observed (L is the dithiolate ligand):

Figure 5. Mechanism for the reaction between ionophore.

Hg2+

3[HgII(L)2] + 2e- h Hg0 + 2[HgII(L)3]-

(7)

2[HgII(L)3]- + 4e- h 2Hg0 + 6L-

(8)

and the

However, in this study, shifts in the potentials of the polarographic waves associated with processes 1 and 2 were observed when 1 was added to a solution of Hg2+ ions containing different supporting electrolyte anions (BF4- and PF6-). This indicates that the supporting electrolyte anion is also involved in the reduction processes, presumably via the BF4- and PF6- anions displacing the ClO4- anion to form the [Hg(SPr)(BF4)] and [Hg(SPr)(PF6)] complexes, respectively (the ClO4- anion is present in all these experiments since the Hg(ClO4)2 salt is added initially). Therefore, a mechanism similar to that shown by eqs 7 and 8 can be written as shown in eqs 9 and 10, which also incorporates the supporting electrolyte anion (X). In this mechanism, eq 9 is associated with reduction process 1 and eq 10 with reduction process 2.

3[HgII(SPr)(X)] + 2e- h Hg0 + [Hg2II(SPr)3(X)3]2- (9) Figure 6. Polarograms for 5 mM Hg(ClO4)2 in methanol containing 0.1 M Bu4NClO4 with excess ionophore: (a) 2-fold and (b) 100-fold molar excess.

to be reduced, but at considerably more negative potentials13 (∼1.7 V vs Ag/Ag+ (0.05 M AgNO3)) than for the processes shown in Figure 6. Therefore, the waves observed by polarography correspond to reduction of the product(s) of a reaction between Hg2+ and the ionophore. To confirm that these processes are not associated with 1b and 1c but rather with 2, a small quantity of 2 was dissolved in methanol, and a polarogram was recorded. It was identical to that shown in Figure 6. When a 100-fold molar excess of the ionophore is added to the solution of Hg2+, the waves labeled as process 1 and process 2 increase in size (Figure 6b). This indicates that the reaction of the ionophore with Hg2+ is not (13) (a) Webster, R. D.; Bond, A. M.; Schmidt, T. J. Chem. Soc., Perkin Trans. 2 1995, 1365. (b) Webster, R. D.; Bond, A. M.; Compton, R. G. J. Phys. Chem. 1996, 100, 1028.

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[Hg2II(SPr)3(X)3]2- + 4e- h 2Hg0 + 3SPr- + 3X- (10)

If 2 is left to dissolve in methanol over 2-3 weeks and then cooled in ice, crystals form which have the empirical formula C15H39S5Cl3O12Hg3 (Table 3, 2a). This is consistent with the structure (SPr)5(HClO4)2ClO4Hg3. The perchloric acid is a biproduct of the initial reaction of 1 with Hg(ClO4)2 and MeOH (Figure 5). Reaction with Pb2+. When equimolar amounts of 1 and Pb(ClO4)2 were mixed together in methanol, white crystals formed after several days. The crystals were not soluble enough in common (deuterated) solvents to enable NMR experiments to be conducted. However, analytical data (Table 3, 3) agreed with the empirical formula C18H18N2O16Cl2Pb. This formula is consistent (14) (a) Benesch, R.; Benesch, R. E.; J. Am. Chem. Soc. 1951, 73, 3391. (b) Hush, N. S.; Oldham, K. B. J. Electroanal. Chem. 1963, 6, 34. (15) Bond, A. M.; Casey, A. T.; Thackeray, J. R. J. Electrochem. Soc. 1973, 120, 1502.

with the structure consisting of two molecules of 1c, one lead atom, and two perchlorate ions. It can be assumed that the initial reaction of Pb2+ with 1 is similar to the reaction of Hg2+ with 1, with cleavage of the carbonyl sulfur bonds. However, the subsequent reaction involves Pb2+ forming a complex with the pyridine fragment of the molecule (1c) and not the propanethiol group, as is the case with Hg2+. Reaction with Cd2+. Cadmium does not interact with 1 in the same manner as Hg2+ and Pb2+. In fact, when equimolar amounts of Cd(ClO4)2 and 1 were mixed in methanol, no new species were formed after several weeks, as shown by thin-layer chromatography. Addition of increasing concentrations of 1 to a 5 mM solution of Cd(ClO4)2 resulted in the polarographic wave for reduction of Cd2+ becoming more drawn out and less reversible in nature. The polarographic behavior is consistent with a weak interaction between 1 and Cd2+, but certainly there is no evidence for the cleavage of the C(O)-S bond in 1 as occurs with Hg2+ and Pb2+. Reaction with Ag+. When equimolar amounts of 1 and AgClO4 were mixed together in methanol, fine white crystals formed after several days. Analytical data (Table 3, 4) agreed with the empirical formula C15H35S5Cl2O8Ag7. This compound is insoluble in all organic solvents tested and most probably exists in a polymeric form of the kind [(PrS)5Ag7(ClO4)2]n. It appears that interaction of Ag+ with 1 occurs in a fashion similar to those of Hg2+ and Pb2+ and leads to a breakdown of the molecule, with subsequent reaction occurring between the silver ion and the thiol reaction product as occurred with Hg2+. As mentioned before, this kind of reaction probably accounts for the decrease in the limiting current seen in the voltammetric studies with silver and for the slight scatter in the potentiometry data. CONCLUSION One aim of this study was to obtain information regarding the complexation behavior between Ag+ and the neutral carrier reagent S,S-dipropyl pyridine-2,6-dicarbothioate, and to relate the excellent response characteristics of a polymer membrane-based sensor made with this reagent1 to this behavior. The study has provided strong evidence using several techniques for the formation of both 1:1 and 1:2 complexes between Ag+ and the ionophore in methanol. The 1:1 complex is the predominant one; however, its formation constant, determined by potentiometry and confirmed by NMR and voltammetry, is low (K1 ≈ 50 L mol-1) compared to the formation constants (102-104 L mol-1) of other cation/ionophore complexes10 used in potentiometric chemical sensors. However, caution should be exercised in making such comparisons, since the stability and chemical studies performed in this work were carried out in methanol and there is no direct experimental evidence that the neutral carrier behaves in exactly the same way in a PVC membrane. This highlights the need for

in situ methods to determine formation constants in the membrane, such as the one reported by Bakker et al.9 as mentioned in the introduction. The high selectivity of the sensor for Ag+ in the presence of most other cations, with the exception of Hg2+, which irreversibly poisons the sensor membrane, implies that the formation constants for complexation between the ionophore and other metal ions are even lower than those for Ag+. The effect of the Hg2+ ion on the sensor membrane has been related to the susceptibility of the ionophore to structurally break down in the presence of this ion. Analogous reactions have been observed before by Satchell and Secemski,16 who studied the hydrolysis of some thiol esters, thiol acids, and thiol anhydrides in aqueous solutions in the presence of metal ions such as Ag+ and Hg2+. They found that metals like Ag+ and Hg2+ which have a strong affinity to bind to sulfur are very effective in promoting the hydrolysis of thiol compounds. In the absence of these ions, hydrolysis reactions are very slow, even in acid conditions. In the present work, the Hg2+ ion is the most reactive of the ions studied and rapidly leads to the cleavage of the C(O)-S bond in the ionophore (1 in Figure 5), with the subsequent precipitation of a product formed between Hg2+ and the thiol fragment. These reactions have been studied in methanol. However, it can be assumed that similar processes can occur at the surface of the polymer membrane sensor containing ionophore 1 in contact with water and the Hg2+ ion. This decomposition reaction is most likely implicated in the strong and irreversible poisoning effect observed when the sensor is placed in a solution containing Hg2+. Similar decomposition reactions were found to occur with Pb2+, which only weakly interferes with the sensor, and Ag+ itself, which poses some difficulties with respect to the functioning of the sensor in silver solutions. The sensor has been found to be highly selective for Ag+ and has excellent response characteristics.1 Thus, the conclusion is drawn that the decomposition reaction of the ionophore, which is fast with mercury, is extremely slow with Ag+ and Pb2+ and hence does not affect the silver response on the short time scale. However, it might be anticipated that the reaction would have severe implications with respect to the lifetime of the sensor, assuming, of course, that the reaction product is not, itself, a selective ionophore for the silver ion. ACKNOWLEDGMENT C.G.G. is grateful for an Australian International Development Assistance Bureau award. R.D.W. thanks La Trobe University for the award of a postgraduate scholarship. We thank the Australian Research Council and the Natural Sciences and Engineering Research Council of Canada for financial assistance.

Received for review January 2, 1997. Accepted May 21, 1997.X AC9700080

(16) (a) Satchell, D. P. N.; Secemski, I. I. Tetrahedron Lett. 1969, 1991. (b) Satchell, D. P. N.; Secemski, I. I. J. Chem. Soc. 1970, 1306.

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Abstract published in Advance ACS Abstracts, July 1, 1997.

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