Detailed Mechanistic Studies into the Reactivities of Thiourea and

Article pubs.acs.org/JPCA

Detailed Mechanistic Studies into the Reactivities of Thiourea and Substituted Thiourea Oxoacids: Decompositions and Hydrolyses of Dioxides in Basic Media Merfat M. Aslabban,† Risikat Ajibola Adigun,† William DeBenedetti,† Wilbes Mbiya,† Morgen Mhike,† Kayode Morakinyo,† Adenike Otoikhian,† Tinashe Ruwona,† and Reuben H. Simoyi*,†,‡ †

Department of Chemistry, Portland State University, Portland, Oregon 97207-0751, United States School of Chemistry and Physics, University of KwaZulu-Natal, Westville Campus, Durban 4014, South Africa

‡

ABSTRACT: Dioxides of methylthiourea (methylaminoiminomethanesulfinic acid, MAIMSA) and dimethylthiourea (dimethylaminoiminomethanesulfinic acid, DMAIMSA) were synthesized and, together with thiourea dioxide (aminoiminomethanesulfinic acid, AIMSA), were studied with respect to their decompositions and hydrolyses in basic aqueous media. All three were stable in acidic media and existed as zwitterions with the positive charge spread out on the 4-electron 3-center N−C−N skeleton and the negative charge delocalized over the two oxygen atoms. All three are characterized by long and weak C−S bonds that are easily cleaved in polar solvents through a nucleophilic attack on the positively disposed carbon center, followed by cleavage of the C−S bond. The sulfur moiety leaving groups are highly unstable, reducing, and rapidly oxidized to S(IV) as hydrogen sulfite in the presence of oxidant. In aerobic conditions, molecular oxygen is a sufficient and efficient oxidant that can oxidize, at diffusion-controlled limits, the highly reducing sulfur species in one-electron steps, thus opening up a cascade of possibly genotoxic reactive oxygen species, commencing with the superoxide anion radical. Radical formation in these decompositions was confirmed by electron paramagnetic resonance techniques. In strongly basic media, decomposition of the dioxides to yield sulfoxylate (SO22−, HSO2−) is irreversible and, in anaerobic environments, will disproportionate to yield more stable sulfur species from HS− to SO42−. Decomposition products were dependent on concentrations of molecular oxygen. Solutions open to the atmosphere, with availability to excess oxygen, gave the urea analogue of the thiourea and sulfate, while in limited oxygen conditions hydrogen sulfite and other mixed oxidation states sulfur oxoanions are obtained. DMAIMSA has the longest C−S bond at 0.188 nm and was the most reactive. MAIMSA, with the shortest at 0.186 nm, was the least reactive. Electrospray ionization−mass spectrometry data managed to detect all of the formerly postulated intermediates.

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INTRODUCTION Our laboratory has, for the past 15 years, been interested in the mechanistic basis of biological activities of thiocarbamides and especially thiourea and its analogues. The thiourea moiety is nearly equivalent to the sulfonamide functionality1−4 in biological and physiological activities.5,6 Several drugs have been constructed that utilize the thiourea backbone either in linear or cyclic form, as the active functionality.7 These include antimalarials8 and novel human immunodeficiency virus (HIV) drugs.9 Their widespread use in drug design necessitates that their mechanism of action be elucidated. Our present research direction now effectively includes the derivation of the relationships between physiological effects and the physical−chemical properties of these biologically active organosulfur compounds. Thiourea derivatives form a vast group of highly reactive and physiologically important compounds.10,11 The substituted thioureas are very toxic.12 Animal studies on the chronic toxicity of thiourea have shown that when it is administered in drinking water, thiourea induces thyroid adenomas and carcinomas in rats.13 No one yet knows the origin of the tumorigenicity of this series of chemicals, but there is speculation that this could be © XXXX American Chemical Society

attributed to thiourea’s strong antithyroid activity which leads to a disruption of the pituitary−thyroid hormonal regulatory system.14 No interaction between thioureas and DNA of thyroid cells has been observed, but, exposure of rats to thiourea is known to result in tumor formation in other organs such as the liver.15 Recent research has reported on phenethyl-5bromopyridyl thiourea as possessing potent anti-HIV activity which when combined with its antioxidant activity, can be an efficient non-nucleoside inhibitor.16 Nearly all relevant sulfur chemistry in the human body is of organic origin. The range of physiological effects associated with organic sulfur chemistry spans from therapeutic to toxic, with several levels of each extreme represented in between.17 There is no other group of compounds that displays such a wide range of biological activity. While the action and effects of organosulfur compounds on human health are welldocumented, the mechanisms by which these effects are Received: April 18, 2014 Revised: September 10, 2014

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dx.doi.org/10.1021/jp5038156 | J. Phys. Chem. A XXXX, XXX, XXX−XXX

The Journal of Physical Chemistry A

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without any further purification. Some batches of DMTU were damp and oily and could not form DMAIMSA according the published method. Instead the product formed slowly turned yellow upon exposure to the atmosphere and at ambient temperatures. The consistency of the DMTU batches was extremely important for the generation of high and consistent yields of DMAIMSA. Several batches were bought but only a few were used for the synthesis of DMAIMSA. Aqueous bromine solutions were prepared by appropriate dilution of liquid bromine followed by a spectrophotometric and iodometric standardization. Standard solutions of sodium hydroxide were purchased from Fisher. Sodium dithionite was purchased from Aldrich Chemical Co (Milwaukee, WI, USA). AIMSA was recrystallized twice from a 50:50 ethanol:water mixture prior to use. Superoxide dismutase (SOD), was sourced from Sigma. The following trap was of the highest purity possible: 5,5-dimethyl-1-pyroline N-oxide (DMPO; Sigma). Doubly distilled deionized water from a Barnstead Sybron Corp. water purification unit capable of producing both distilled and deionized water (Nanopure) was used for the preparation of all stock solutions and their subsequent dilution. All solutions were prepared fresh unless otherwise stated. Adventitious metal ions are very disruptive to reactions involving sulfur compounds. We utilized inductively-coupled plasma mass spectrometry (ICPMS) to quantitatively evaluate the concentrations of a number of metal ions in the water used for our reaction medium. ICPMS analysis showed negligible concentrations of iron, copper, and silver and approximately 1.5 ppb cadmium and 0.43 ppb lead. The use of chelators to sequester metal ions gave kinetics and reaction dynamics indistinguishable from those run in deionized water. Synthesis of MAIMSA. Methylaminoiminomethanesulfinic acid was prepared according to a modified synthetic procedure of that used for the preparation of DMAIMSA.28 Briefly, 9.04 g of MTU (0.10 mol) was dissolved in 80 mL of a 50% acetonitrile solution which was chilled in an acetone/CO2 ice bath at −59 °C. To this ice-cold solution, 2 equiv (0.2 mol) of hydrogen peroxide, which was measured as a 20.41 mL aliquot of a 30% concentrated solution, was added dropwise with the rate of addition maintained such that the temperature of the reaction mixture did not exceed −39 °C. The 2:1 ratio of hydrogen peroxide to MTU had to be strictly maintained in order to avoid the production of mixtures of oxoacids. Addition of even slightly more H2O2 than the 2:1 ratio gave product mixtures containing MAIMSA, the sulfonic acid, MAMSOA, and sulfate. The frozen mixture was allowed to sit until it melted out and then stirred at room temperature for 2 h. The resulting colorless needlelike crystals were filtered and washed twice with 50% acetonitrile solution and deionized water and then dried in a desiccator. Typical yields were approximately 65%. Further recrystallization and purification was still performed using the same strength acetonitrile solution. Characterization of MAIMSA. MAIMSA was characterized by X-ray crystallography and titrimetry. It is expected that the oxidation state of the sulfur center in MAIMSA should be +2 and that the sulfur center requires 4 electrons for a full oxidative saturation to the +6 sulfate oxidation state. A titration of MAIMSA vs aqueous bromine (enhanced by excess iodide and soluble starch) showed that 2 mol of bromine was required for a full oxidation of the putative MAIMSA synthesis product. The 2:1 ratio concluded that our synthesized product was MAIMSA.

expressed are not known. Drug design and ability to predict physiological effects is dependent upon our ability to understand the mechanism of metabolic activation of the relevant compound. In the physiological environment, due to the nucleophilicity of the sulfur atom, all bioactivation reactions of thioureas and thiocarbamides are oxidative.18−21 All physiological activities observed from thioureas are derived from their oxidation metabolites. This assertion is supported by the fact that genotoxicity of thioureas is correlated to the degree of desulfurization during bioactivation.22 Desulfurization would suggest oxidation of the sulfur center up until it is oxidatively saturated to +6 oxidation state where the C−S bond cleaves to give sulfate. Any cleavage of the C−S bond before the sulfur atom is oxidatively saturated should produce reactive and subsequently genotoxic metabolites. Previous work from our laboratory had concentrated on the decomposition mechanisms of rongalite23−26 and aminoiminomethane sulfinic acid (AIMSA; Chart 1),27 both industrially Chart 1. Structures of the sulfinic acids studied in this work: AIMSA (top left), MAIMSA (top right), and DMAIMSA (bottom).

important compounds. AIMSA, the dioxide of thiourea, is a well-known commercial compound and available from chemical vendors. It is used in the textile industry for the generation of free radicals to initiate polymerizations. We decided, in this work, to determine the reactivities of the first stable metabolites of thioureas, the dioxide analogues (sulfinic acids). For a complete study, we utilized dioxides of thiourea, methylthiourea (MTU) and dimethylthiourea (DMTU). The dioxide for MTU, methylaminoiminomethanesulfinic acid (MAIMSA), had been synthesized and isolated in our laboratories before but had never been characterized. We synthesized and characterized it for this work (vide inf ra). The dioxide for DMTU, dimethylaminoiminomethanesulfinic acid (DMAIMSA), had been prepared and isolated in our laboratory28 and was resynthesized by the same method for the work reported here. This work involves new techniques and approaches that were not available in our previous work of 15 years ago (electrospray ionization−mass spectrometry (ESI-MS), electron paramagnetic resonance (EPR), and diode array capabilities). We also now have been able to conclusively characterize the MAIMSA via X-ray crystallography.

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EXPERIMENTAL SECTION DMTU, MTU, thiourea, hydrogen peroxide, and acetonitrile (Sigma-Aldrich, St. Louis, MO, USA) were purchased and used B



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a NesLab RTE-101 thermostat. Mass spectra of product solutions were taken on a Thermo Scientific LTQ-Orbitrap Discovery mass spectrometer (San Jose, CA, USA) equipped with an electrospray ionization source operated in the negative mode. All EPR spectra were recorded on a Bruker Biospin e-scan spectrometer designed to perform EPR measurements in the X-band range at room temperature. Oxygen Measurements. Oxygen uptake measurements were done at room temperature on a Gilson Oxy 5/6 oxygraph. A YSI 5331 oxygen probe was inserted into a stirred and thermostatically controlled reaction vessel coupled to an amplifier and recorder. The oxygen probe was a complete polarographic system consisting of a platinum cathode, silver anode, and KCl solution held captive around the electrodes by a Teflon membrane fastened with an O-ring. The response time is 90% in 10 s, and it takes 30 s to reach steady state mainly because of the membrane. Since the probe is a complete system by itself, it is relatively unaffected by, and does not offset, its external environment. Adequate stirring eliminated reading error due to oxygen depletion in the vicinity of the membrane.

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RESULTS The observed reaction dynamics under all constraints featured here were effectively identical for AIMSA, MAIMSA, and DMAIMSA (Figure 1). DMAIMSA reactions, under identical conditions, were the most rapid, followed by AIMSA with MAIMSA reactions being the most sluggish. The dioxides all had absorption peaks in the UV region with λmax values of 270, 257, and 263 nm for AIMSA, MAIMSA, and DMAIMSA, respectively. MAIMSA shows a variation in this peak based on the pH. In water and pH 6, λmax = 263 nm. In basic environments, pH > 9.4, this peak shifts to the observed 257 nm. Without an oxidant, the dioxide solutions are stable in acidic conditions, pH < 3, with halflives of days, even in the presence of acetate buffers. In aerobic conditions and pH conditions higher than 6.5, the dioxides decompose slowly to give the corresponding urea analogue and hydrogen sulfite, HSO3−. The reactions shows a monotonic decrease in the 270−257 nm peak and the emergence of a species that absorbs strongly at 315 nm after a delay.

Figure 1. (A) Multiple scans taken every 30 s for the decomposition of AIMSA (0.001 M) in 0.05 M NaOH. The 315 nm peak, due to dithionite, S2O42−, increases monotonically after a delay. The 270 nm peak, due to AIMSA, initially decreases until there is contribution from dithionite which has significant absorbance at 270 nm. (B) Scans taken every 30 s of a 0.001 M MAIMSA solution with 0.05 M NaOH. The absorbance maximum of MAIMSA shifts from 263 to 257 nm in basic environments. There is less contribution from dithionite at 257 nm compared to 270 nm.

Me(H)N(H 2N)CSO2 H + 2Br2(aq) + 3H 2O → Me(H)N(H 2N)CO + 6H+ + SO4 2 − + 4Br − (R1)

The crystal structure showed a C2/c space group with two important features: (a) the C−S bond was inordinately long, at 0.1860 nm. This can be compared with the expected bond length of 0.179 nm if one simply added the covalent radii of the two bonding atoms; and (b) the two S−O bonds are equivalent, indicating a delocalization of the negative charge over the two oxygen atoms, and that the C−N bond lengths are between the CN double bond and C−N single bond lengths, suggesting a delocalization of the positive charge in the zwitterionic form of MAIMSA. Instrumentation. Absorptivity coefficients were measured on a Perkin−Elmer Lambda 25S UV−vis spectrophotometer. For faster reactions, a Hi-Tech Scientific SF-DX2 stopped-flow spectrophotometer was used to follow dithionite formation at 315 nm (εmax = 8043 M−1 cm−1). KinetAsyst 2.1 software was used for data acquisition and analysis. Temperature control was maintained with

Figure 2. Induction period followed by zero-order formation of dithionite. [NaOH] = 0.5 M; [AIMSA] = (a) 1.0 × 10−3, (b) 1.25 × 10−3, (c) 1.20 × 10−3, (d) 1.75 × 10−3, and (e) 2.0 × 10 −3 M. C



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The rate of formation of dithionite and the amount of dithionite formed, but not the induction period, are catalyzed by buffer anions. Figure 4 shows a series of experiments in which the pH is controlled either strictly by flooding with NaOH or by the use of a Britton−Robinson buffer system. A marked increase in the rate of formation of dithionite is observed with the corresponding buffered reaction solutions. The reaction also displayed a mild primary salt effect with the reaction being slightly faster at higher ionic strength. There was no linear relationship between the inverse of the induction time or the rate of formation of dithionite on the inverse of the ionic strength, indicating that the reaction is complex and is not controlled by a single rate-determining step. Such a relationship, should it exist, is a hallmark of the primary salt effect. Oxygen Effects. Our previous work had established the pivotal role of oxygen in the decomposition of thiourea dioxides.32 Oxygen controlled both the induction period and

The species that absorbs at 315 nm has been established to be dithionite, S2O42−. Its nature has been established as well as its absorptivity coefficient at 315 nm. Formation of dithionite, after the induction period, appears to be essentially zero order (Figure 2). Zero-order kinetics are prevalent in surface kinetics, sublimations, and reactions in which there is a time lag between the reaction that forms the precursor species and the reaction that utilizes this precursor species to form the observable. The dithionite is later autooxidized to hydrogen sulfite at a much slower rate.29 If the reaction is observed for a longer period, depletion of the dithionite will be consumed through autoxidation.30,31 S2 O4 2 − + 1/2 O2 + H 2O → 2HSO3−

(R2)

The observation of dithionite consumption can be accelerated by the addition of an oxidizing agent such as H2O2 (see Figure 3).

Figure 3. Prolonged observation of the decomposition of DMAIMSA in basic medium and in the presence of an oxidant. Dithionite is further oxidized to sulfate through hydrogen sulfite. [DMAIMSA]0 = 1.0 × 10−3 M; [NaOH]0 = 0.50 M. These solutions were purged with argon for 3 min. [H2O2]0 = (a) 2.2 × 10−3, (b) 4.4 × 10−3, (c) 8.8 × 10−3, (d) 1.1 × 10−2, (e) 2.2 × 10−2, and (f) 4.4 × 10−2 M.

Figure 5. Effect of oxygen on decomposition of MAIMSA: (a) under normal ambient oxygen conditions; (b) after purging reaction solution with argon for 6 min; (c) after purging for 10 min. The same purging done on DMAIMSA gives similar dynamics but at a much more rapid rate (see Figure 6). [MAIMSA]0 = 0.001 M.

Figure 4. Comparison of reaction solutions in which pH is maintained by buffer and those maintained strictly by NaOH. [AIMSA]0 = 0.001 M. D



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Observations at 270 nm. Depletions of the dioxides were all observed at 270 nm, even though this peak shifted with prevailing pH for the substituted dioxides. The corresponding urea products also absorb at 270 nm but with lower absorptivity coefficients. Figure 8 shows kinetics traces of the

Figure 6. Effect of oxygen on the alkaline decomposition of DMAIMSA: (a) ambient conditions; (b) after purging the reaction solution with argon for 6 min; (c) after 10 min purging. [DMAIMSA]0 = 0.001 M.

the amount of maximum dithionite formed before its final autoxidation to hydrogen sulfite. Figure 5 shows the strong influence of oxygen on decomposition of alkaline MAIMSA. Completely Anaerobic Conditions. Decomposition reactions of AIMSA, MAIMSA, and DMAIMSA were run in strictly anaerobic conditions in a glovebox. All reaction solutions used for the reactions were purged in argon while in the glovebox, and reactions were commenced when oxygen concentrations had gone below 1 ppb. The stopped-flow sample handling unit was housed inside the glovebox. The kinetics traces were obtained at both 270 and 315 nm and superimposed on those obtained in aerobic environments. Figure 7 shows these four experiments. In both

Figure 8. Depletion of AIMSA monitored at 270 nm. Higher pH conditions reduce the induction period and rapidly produce dithionite which contributes strongly to the absorbance observed at 270 nm. [AIMSA] = 0.001 M; [NaOH]0 = (a) 0.10, (b) 0.20, (c) 0.25, (d) 0.30, (e) 0.40, and (f) 0.50 M.

decomposition of AIMSA at 270 nm at ambient conditions. These data erroneously predict that NaOH inhibits decomposition of AIMSA. Higher pH conditions reduce the induction period, thus asserting an earlier production of dithionite. Dithionite has a substantial absorptivity at 270 nm, thus erroneously predicting inhibition by NaOH. All reactions were run in capped reaction cuvettes. Oxygen available for the reaction is limited to that initially contained in the cuvette. When reactions are run in open cuvettes, complex dynamics are obtained (see Figure 9). With open cuvettes, reaction kinetics become irreproducible and stochastic. Though every effort was made to run traces shown in Figure 9 at exactly the same conditions, even a slightly greater agitation of one solution over another would deliver completely different kinetics traces. All reactions shown in Figure 9 were not stirred. Kinetics Observations at Both 270 and 315 nm. Using a diode array spectrophotometer, simultaneous kinetics observations were made at both 270 and 315 nm. Figure 10 shows a reconstruction of the absorbance data at both wavelengths with vertical lines connecting similar events in the corresponding traces of the same initial conditions. As expected, the observed increase in absorbance at 270 nm coincides with the end of the induction period and commencement of dithionite production. Absorbance at 270 nm commences to decrease again when dithionite reaches its peak absorbance. Radical Species Generation. EPR spectra (Figure 11) were taken as the reaction proceeded, using DMPO as the trap. Initial spectra were complex and appeared to involve a superposition of two or more radical species. An attempt to simulate the expected spectra on the basis of two species, superoxide anion radical and sulfite anion radical, failed, and the program could not pick up and separate the expected peaks and their hyperfine coupling constants.

Figure 7. Aerobic and completely anaerobic conditions. No dithionite is formed in the absence of oxygen, but decomposition of DMAIMSA occurs in both conditions: (a) aerobic decomposition of DMAIMSA; (b) formation of dithionite; (c) anaerobic decomposition of DMAIMSA; (d) no dithionite formed in anaerobic condition. [DMAIMSA]0 = 0.001 M.

aerobic and anaerobic conditions, decomposition of DMAIMSA is observed at nearly the same rate as in aerobic environments, but, in anaerobic environments, no dithionite is formed. E



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Figure 9. Absorbance traces in open cuvettes at 270 nm while varying NaOH from (a) 0.1 to (e) 0.5 M. Reaction dynamics are affected even by the slightest agitation of the reaction cuvettes. Overall, reaction kinetics is no longer reproducible. [AIMSA]0 = 0.001 M.

At the point where dithionite reaches its maximum concentration, a singlet radical peak, not adducted to DMPO, shows up at 3475 G. This is a well-known radical, the sulfoxyl anion radical, SO2•−. A two-dimensional (2-D) scan was made which showed that this was the only radical species present (Figure 12). Another series of experiments were undertaken to prove that these radicals are derived solely from the dioxides (Figure 13). ESI-MS Experiments and Data. All of the dioxides showed the same types of speciations, except for the rate at which the speciation species are attained. In this work we will only present data derived from MAIMSA and DMAIMSA decompositions. DMAIMSA is utilized because it decomposes at the fastest rate and produces most of the metabolites within 2 min of reaction, while MAIMSA is used because it is the most stable and can be monitored at the beginning of the reaction before decomposition has significantly proceeded. Figure 14A shows the mass spectrum for MAIMSA in water, at a pH of 6.0. Even at this slightly acidic pH, before addition of NaOH, the spectrum shows a strong peak for the sodium salt of MAIMSA at m/z = 145.0 with peaks for protonated thiosulfate, S2O32−, at 111.05 m/z and for hydrogen sulfite, HSO3−, at 81.05. At these pH conditions, MAIMSA is mostly undissociated and shows a strong peak at m/z = 123.02. There are other unidentified peaks at 186.2 and 267.02. Figure 14B shows the spectrum from DMAIMSA 2 min after addition of 0.05 M NaOH. As expected, the sodium salt peak at 145.0 m/z vanishes and we see the emergence of the dithionite peak at 129.05 m/z. Since this was run in aerobic conditions, autoxidation occurs and sulfate begins to show as the hydrogen sulfite at 97.04 m/z. This bisulfate peak increases at the expense of the hydrogen sulfite peak, which essentially vanishes. The thiosulfate peak at 111.05 m/z increases. The expected full oxidation product, methylurea, begins to show up at 90.97. Further incubation of this solution in an open cuvette shows an increase in the dimethylurea and hydrogen sulfite peaks. The dithionite peak lingers for extended periods of up to 8 h. It takes over 48 h for a complete consumption of the dithionite, and in the absence of an oxidant, the major oxidation products are dimethylurea and surprisingly, thiosulfate, not hydrogen sulfite. In the presence of any oxidant, the thiosulfate is rapidly converted to sulfate. In a closed cuvette, where oxygen is limited, dithionite and thiosulfate retain their levels indefinitely.

Figure 10. Absorbance traces at 270 and 315 nm. The 270 nm peak estimates the concentration of AIMSA and the 315 nm peak follows the dithionite concentrations. [NaOH]0 = 0.50 M; [AIMSA]0 = (a) 0.00225, (b) 0.002, (c) 0.0175, and (d) 0.0015 M.

Figure 11. EPR spectra of the sulfoxyl anion radical at 3475 G on an X-band EPR. Sizes of singlet peaks are determined by the initial DMAIMSA concentrations in excess NaOH. [NaOH]0 = 0.50 M; [DMAIMSA]0 = (a) 0, (b) 0.005, (c) 0.001, (d) 0.002, (e) 0.003, (f) 0.004, and (g) 0.005 M. F



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Figure 12. Two-dimensional scan for radicals in the decomposition of MAIMSA (0.001 M) in 0.05 M NaOH. Spectrum was scanned between 200 and 800 s. No other radical species are observed at the end of the reaction.

Figure 13. No radicals are obtained in the absence of AIMSA. Only if dithionite is produced are sulfoxyl radicals formed: (A) formation of dithionite in the reaction of aerobic AIMSA decomposition in 1 × 10−3 M AIMSA air-saturated solution ([NaOH]o = (a) 0.05 and (b) 0.10 M); (B) electron paramagnetic resonance spectroscopy of 1 × 10−3 M AIMSA and INaCl = 1 M ((pure) no AIMSA; (b) with 0.05 M NaOH; (c) with 0.1 M NaOH).

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MECHANISM Table 1 shows the crystal structures of all the three thiourea dioxides. In solid state, they all exist in the zwitterionic forms with the negative charge delocalized over the two oxygen atoms and a 4-electrom 3-center framework at the N−C−N site. In all of he dioxides, the C−S bond is elongated and, hence, is easily cleaved. Reactivity is related to the length of the C−S bond. At 0.188 nm,

Dithionite is a powerful reducing agent and very easily oxidized.25,30,31 It undergoes rapid disproportionation in aqueous solution to form more stable sulfur compounds in different oxidation states ranging from HS− to SO42−.33 The main products of the decomposition are sulfite and thiosulfate:34 2S2 O4 2 − + H 2O → 2HSO3− + S2 O32 −

(R3) G



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Figure 14. (A) ESI spectrum of MAIMSA in water, pH = 6.0. It shows a strong peak for the sodium salt of MAIMSA, Me(H)N(NH)CSO2Na at m/z = 145.00. The dissociated salt, the sulfinate, shows up at 123.02; HSO3−, at 81.04; and thiosulfate, at 111.05. (B) ESI spectrum of a 0.05 M NaOH solution of DMAIMSA after 2 min of incubation. The thiosulfate peak becomes more pronounced as well as the HSO4− peak at 97.04. Protonated dithionite gives a strong peak at 129.05. Product dimethylurea shows up at 90.98.

the C−S bond in DMAIMSA is the longest, explaining why DMAIMSA is the most reactive. The carbon center is positively charged and is a prime position for a nucleophilic attack. In the

presence of water, Scheme 1 represents the plausible pathway toward the cleavage of the C−S bond to deliver an oxidatively unsaturated S atom at an oxidation state of +2: H



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DMTU is one of the most efficient hydroxyl radical scavengers.36 It efficiently scavenges toxic oxygen metabolites in vitro and reduces oxidative injury in many biological systems.37,38 A lot of research has been done on the biological effects of DMTU, but none of the research has been able to give us a predictive tool for toxicity of DMTU. Although in most situations it is effective in decreasing oxidant mediated injury, it has, however, inexplicably failed, on occasion, to reduce injury in some biological systems where oxygen metabolites were ostensibly causing damage.38,39 It has also been experimentally shown that increasing the DMTU dose did not increase protection and may, in fact, be associated with more injury.40 This formally unrecognized toxicity of DMTU provides a possible explanation for the apparently conflicting reports regarding the efficacy of DMTU in cytoprotectivity. Our work here helps to clarify this ambiguity. Though it initially efficiently scavenges toxic oxygen species, the oxidative DMTU metabolites generated (such as the dioxides studied here in this work) are also equally reactive and generate, on their own, toxic metabolites as well. The physiological pH, being basic, is conducive to the decomposition of these dioxides and other oxoacids. Stable metabolites, hydroxylated, can be eluted through the kidneys after they have quenched toxic oxygen metabolites. DMAIMSA’s high reactivity in the basic environment militates against this and instead produces a cascade of further toxic metabolites. Ultimately, the occurrence of increased stomach size, lethargy, and other behavioral changes in rats provides support to the premise that DMTU is, intrinsically, toxic.38

Table 1. X-ray Structural Data for Some Synthesized Thiourea-Based Metabolites property

AIMSA

MAIMSA

DMAIMSA

C−S bond length (Å) S−O bond length (Å)

1.867(7)

1.860(3)

1.880(2)

1.496(3)

1.472(1)

1.476(2)

C−N bond length (Å)

1.296(5)

1.499(2) 1.307(2)

1.479(2) 1.303(3)

both equivalent 1.68 SO2C(NH2)2

1.293(2) 1.449 SO2C(NHMe)NH2.H2O

1.3043 1.496 SO2C(NHMe)2

fast

fast

fast

specific gravity empirical formula rate of oxidation

Scheme 1. Initial Hydrolysis

Scheme 1 uses water as the nucleophile. The hydrolysis reaction is essentially irreversible. Buffer anions and the hydroxide anion would be much more efficient nucleophiles than water. This explains the higher rate of hydrolysis in basic media and in high pH buffers. Our previous studies had established the sulfoxyl anion radical as the sole precursor to the formation of dithionite.23 Its rate of reaction with oxygen is known to be essentially diffusion controlled, yielding a superoxide anion radical: HSO2− + O2 → SO2•− + O2•− + H+

(R4)

2SO2•− ⇄ S2 O4 2 −

(R5)

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*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was funded partly by a grant from the Saudi Arabian Culture Department to M.M.A. and Research Grant Number CHE 1056311 from the National Science Foundation and a partial research professorship allocation from the University of KwaZulu-Natal given to R.H.S.

Reaction R5, formation of dithionite, competes with further oxidation of the sulfite anion radical to hydrogen sulfite: SO2•− + O2 + H 2O ⇄ HSO3− + O2•− + H+

AUTHOR INFORMATION

Corresponding Author

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(R6)

REFERENCES

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Since reaction R4 is diffusion controlled, the amount of dithionite formed is determined by the available oxygen. Hence, in strictly anaerobic conditions, no dithionite formation is observed (see Figures 7 and 13).

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CONCLUSION Our EPR experimental data could not conclusively isolate the different radical species that could be formed in this mechanism. A previous study from our laboratory on rongalite had found a relatively clean spectrum of the sulfite anion-radical adduct with DMPO.23 The sulfite anion radical was expected in the decomposition mechanism of rongalite. The EPR spectrum for this adduct is close to the known 1:2:2:1 splitting for the DMPO superoxide anion radical with a hyperfine coupling constant of aN = 12.99 G.35 With the expected hydroxyl ion radical and a possible cascade that includes S2O5.‑, one would expect a complex and illegible EPR spectrum as was obtained in our experiments. I



Article

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