Decomposition of N-chloroglycine in alkaline aqueous solution: kinetics and mechanism.

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The Decomposition of N-Chloroglycine in Alkaline Aqueous Solution: Kinetics and Mechanism Mária Szabó, Zsolt Baranyai, László Somsák, and Istvan Fabian Chem. Res. Toxicol., Just Accepted Manuscript • DOI: 10.1021/acs.chemrestox.5b00084 • Publication Date (Web): 07 Apr 2015 Downloaded from http://pubs.acs.org on April 11, 2015

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Chemical Research in Toxicology

The Decomposition of N-Chloroglycine in Alkaline Aqueous Solution: Kinetics and Mechanism Mária Szabó#, Zsolt Baranyai#, László Somsák§, István Fábián*,# #

Department of Inorganic and Analytical Chemistry, University of Debrecen, Debrecen, Hungary §

Department of Organic Chemistry, University of Debrecen, Debrecen, Hungary

KEYWORDS: N-chloroglycine, N-formylglycine, N-oxalylglycine, glyoxylate, decomposition kinetics, redox mechanism

ABSTRACT:

The decomposition kinetics and mechanism of N-chloroglycine (MCG) was

studied under very alkaline conditions ([OH−] = 0.01 – 0.10 M). The absorbance change is consistent with two consecutive first-order processes in the 220 – 350 nm wavelength range . The first reaction is linearly dependent on [OH−] and interpreted by the formation of a carbanion from MCG in an equilibrium step (KOH) and a subsequent loss of chloride ion from this intermediate : kobs1 = KOH k1 = (6.4 ± 0.1) ×10-2 M−1s−1, I = 1.0 M (NaClO4), T = 25.0 °C. The second process is assigned to the first-order decomposition of N-oxalylglycine which is also formed as an intermediate in this system, kobs2 = (1.2 ± 0.1) × 10-3 s−1. Systematic 1H and

13

C

NMR measurements were performed in order to identify and follow the concentration changes of the reactant, intermediate and product. It is confirmed that the decomposition proceeds via the 1 ACS Paragon Plus Environment


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formation of glyoxylate ion and produces N-formylglycine as a final product. This compound is stable for an extended period of time, but eventually hydrolyses into formate and glycinate ions. A detailed mechanism is postulated which resolves the controversies found in earlier literature results.

INTRODUCTION In recent years, physiological processes and environmental relevance have generated immense interest in the redox reactions of chlorine and hypochlorous acid with amines, amino acids, peptides and proteins.1-19 These reactions lead to the formation of various N-chlorinated amines under various conditions.20, 21 Hypochlorous acid is formed in vivo by the oxidation of chloride ion with hydrogen peroxide catalyzed by the myeloperoxidase enzyme.22, 23 The formation of HOCl has an essential role in destroying various pathogens in living systems.15,

24, 25

The

consequences of the formation of N-chloramines in biological systems are controversial. These species are secondary disinfecting agents but also contribute to adverse effects in living cells.26-35 Penetration of these molecules into the cell is a key issue. Trans-chlorination, i.e. the transfer of chlorine from an N-chloramine to another amino derivative, has outstanding importance in this

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phenomenon because the penetration ability of different species may be significantly different.36 The metabolites formed from N-chlorinated compounds may also have strong impact on biological processes. The very same species also play important roles in water treatment technologies.37 Quite often they are formed spontaneously from dissolved organic compounds upon dosing chlorine or hypochlorous acid to water.

When the chlorinating agent is present in excess, breakpoint

chlorination occurs, i.e. ammonia is completely oxidized leading to the formation of N2. It was shown earlier that glycine may have a profound effect on this process.38

Under certain

conditions N-chloramines may exist in water for an extended period of time. They exhibit disinfecting activities and sufficiently kill microorganisms, though their efficiency is generally less than that of chlorine.39 On the other hand, they clearly present a health risk, thus, their removal from the finished water is necessary. Earlier studies on the formation of N-chloramines from amino acids revealed that the main reaction path occurs between the deprotonated amino acid and HOCl.23, 40-43 The second order rate constants for this reaction are within the range of 107 – 108 M−1s−1. The calculated pH profile of the pH-dependent rate constants for the formation of N-chloroglycine (monochloroglycine, MCG) goes through a maximum at around pH 8.5 (Figure S1).41 Systematic studies on the kinetics and mechanism of the decomposition of N-chlorinated amino acids postulate the formation of imines which undergo further reactions.20, 21, 23, 44-50 Some of these investigations suggests that decarboxylation and the loss of chloride ion occurs via a concerted Grob fragmentation mechanism.51-53 Earlier results on the decomposition kinetics are somewhat controversial. In a fairly detailed study on the decomposition of MCG, Hand and coworkers reported the formation of formaldehyde, a small amount of glyoxylic acid and



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ammonia in a pH independent (neutral – slightly alkaline pH range) first-order process.45 A slight deviation from the first order behavior was observed at the beginning of the kinetic traces. This effect was attributed to the decomposition of a N,N-dichloroamino acid which was supposedly formed in small amounts in this system.

Other studies also reported that the

decomposition rate is independent of pH in alkaline solution.47, 49 However, according to our experience the stability of the MCG stock solution is pH-dependent in the entire alkaline pH range (pH = 7.0 – 13.0) with the highest stability at pH ~ 7.0, and Armesto et al. reported linear dependence of the reaction rate on the hydroxide ion concentration at high pH.48 The existence of the dichloro species is also questionable in MCG stock solutions when glycine is used in excess. The reaction of glycine with HOCl is very fast, and HOCl is most likely totally consumed before it could react with MCG.41 All of these studies imply that the main product of the decomposition of MCG is fomaldehyde. The main objective of our present investigation is to resolve some of the contradictions regarding the decomposition kinetics and mechanism of MCG in alkaline solution. It will be shown that the reaction is far more complex than assumed before and the decomposition products may have a profound effect on experimental work with MCG. Identification of the decomposition products and their formation may also prove to be significant in the interpretation of the exact role of MCG in biological systems. MATERIALS AND METHODS Materials. Chloride ion free sodium hypochlorite was prepared as described earlier.54, 55 The stock solutions of NaOCl were stored at 5 ºC in the dark and were standardized before use. After acidification, KI was added to an aliquot of the solution and the iodine formed was titrated by standardized Na2S2O3 solution yielding the concentration of hypochlorite ion.56 Another portion


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of the stock solution was titrated with HClO4 solution. In this titration, two equivalence points were observed allowing the calculation of the concentration of OCl− and the excess NaOH. An excellent agreement was found in the concentration of hypochlorite ion from the two analytical methods. Glycine (Sigma-Aldrich), and sodium glyoxalate (Sigma-Aldrich) were of reagent grade quality and used without further purification. In the experiments with formaldehyde, sufficient amount of paraformaldehyde (Sigma-Aldrich) was dissolved in the reaction mixture. Formylglycine was prepared as described earlier.57

N-

The samples were prepared in doubly

deionized and ultrafiltered water from a MILLI-Q RG (Millipore) water purification system which was distilled prior use. The kinetic measurements were made at 25 ± 0.1 °C and 1.0 M ionic strength set with NaClO4 prepared from HClO4 (Reanal) and Na2CO3 (Reanal) as described earlier.58 The hydroxide ion concentration was adjusted with NaOH solution. Instruments and methods. Iodometric and pH titrations were made with a Metrohm 721 NET Titrino system equipped with Metrohm 6.0451.100 combination platinum and Metrohm 6.0262.100 combination glass electrodes, respectively. Spectrophotometric measurements were made with a Hewlett-Packard 8543 UV/VIS diode array spectrophotometer (equipped with a built-in magnetic stirrer). The possibility of unwanted photoreactions was tested by using different illumination protocols.59

Photochemical

interference was not observed in this system. Some of the experiments were performed in a Shimadzu UV-1800 double beam spectrophotometer. Measurements were made in stoppered tandem quartz cuvettes of 8.75 mm light path. The temperature of the cell was controlled by a built-in thermoelectric Peltier device in both instruments. In most cases, the baseline of the



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spectrophotometric experiments was recorded with the excess glycine, thus, spectral contribution of this species in the far UV region (below 240 nm) was eliminated. Kinetic experiments were performed by mixing alkaline glycine and hypochlorite ion solutions and monitoring the spectral changes over the 220 – 350 nm spectral range. The calculations were made with MatLab60 and data fitting was made with the program package ORIGIN61 using non-linear least-squares routines. 1D (1H and 13C) and 2D (COSY, NOESY, HSQC and HMBC) NMR measurements were made by using a Bruker DRX 400 (9.4 T) NMR spectrometer equipped with a Bruker VT-1000 thermo-controller and BB inverse z gradient probe (5 mm). With a few exceptions, each solution was prepared in H2O and DSS (4,4-dimethyl-4-silapentane-1-sulfonic acid) in D2O was added to the sample in a capillary as an external standard for 1H (0 ppm). The 1H-NMR spectra were recorded by using the standard Bruker watergate pulse sequence for the suppression of water proton signal. In each 1H-NMR experiment, 24 scans were collected with 16K data points using a sweep width of 5995 Hz, a pulse angle of 90°, an acquisition time of 1.366 s and relaxation delay of 1 s.

13

C-NMR spectra were recorded in J-modulated decoupling mode with the

following parameters: 12000 scans, 32K data points, a sweep width of 22075 Hz, a pulse angle of 90°, an acquisition time of 0.74 s and relaxation delay of 5 s. The COSY, NOESY, HSQC and HMBC spectra were collected by using gradient pulses in the z direction with the standard Bruker pulse programs. For NOESY spectra the mixing time (D8) was 300 ms. The spectra were analyzed with the Bruker WinNMR software package. A MicroTOF-Q type Qq-TOF MS instrument (Bruker Daltonik, Bremen, Germany) was used for the MS measurements with an ESI source in negative ion mode. The spray voltage was 4 kV and the temperature of the drying gas (N2) was kept at 180 °C. The spectra were accumulated


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and recorded with a digitizer at 2 GHz sampling rate. The mass spectra were calibrated using the exact masses of the clusters generated from the electrosprayed solution of sodium trifluoroacetate (NaTFA). The spectra were analyzed with the DataAnalysis 3.4 software from Bruker. RESULTS AND DISCUSSION Preliminary results. The formation of MCG from hypochlorite ion and glycine is very fast under the conditions applied and should be considered complet before the decomposition reaction starts for all practical purposes. The characteristic absorbance band of MCG with λmax = 255 nm was observed immediately after mixing the reactants and the subsequent absorbance decay is due to the decomposition of this species (Figure 1).

Figure 1. The immediate formation and subsequent decay of MCG (λmax = 255 nm) upon mixing aqueous solutions of hypochlorite ion and glycine in excess. The spectra were recorded in 10 s intervals up to 5000 s (for sake of simplicity only every fourth spectrum is shown). CGLY0 = 2.00 × 10-4 M, CMCG0 = 1.00 × 10-3 M, COH− = 0.054 M, I = 1.0 M (NaClO4), T = 25.0 °C. Note that glycine consumes OCl− immediately upon mixing the reactants and MCG forms in equivalent concentration of OCl− added.



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The kinetic measurements were performed by using different hypochlorite ion concentrations and glycine in excess. The initial absorbance is a linear function of OCl− concentration at 255 nm and no distortion of the characteristic absorbance band at λmax = 255 nm was observed. When an excess of OCl− was used, a clear sign of the formation of another species, presumably N,N-dichloroglycine was observed in the spectra (Figure S2). These results confirm that only the formation of MCG need to be considered under the conditions applied here. NH2CH2COO− + OCl− = ClNHCH2COO− + OH−

(1)

Kinetics. In the kinetic experiments, MCG was prepared by mixing alkaline solutions of hypochlorite ion and glycine at least in 20 % excess, and the subsequent spectral changes were followed. Typical kinetic traces are shown in Figure 2.


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Figure 2. Typical kinetic traces for the decomposition of MCG at 280 nm (a) and 228 nm (b) at different concentrations of excess glycine. The solid lines are fitted traces to equations 2 and 3. CMCG0 = 3.00×10−3 M, COH− = 0.054 M and CGLY0 = 1.50×10−3 (◊), 3.00×10−3 (□), 1.20×10−2 (○), 2.70×10−2 (∆); I = 1.0 M (NaClO4), T = 25.0 °C. Note that glycine consumes OCl− immediately upon mixing the reactants and MCG forms in equivalent concentration of OCl− added. At 280 nm simple first-order behavior was observed and the traces could be fitted to a single exponential function according to equation 2.

A = A0e−k obs1t + A∞

(2)

where A is absorbance and kobs1 is the pseudo-first order rate constant for the decomposition of MCG. The deviation of A∞ from zero at the highest glycine concentration (Figure 2a) confirms that at least one of the products has a slight contribution to the absorbance at 280 nm. At lower wavelengths, systematic deviations were observed between the measured and fitted first-order kinetic curves. Below 240 nm, the kinetic profiles are consistent with the formation and disappearance of an intermediate (cf. the kinetic traces at 228 nm, Figure 2b). Singular value decomposition analysis of the time resolved spectral data were performed for individual kinetic runs at a given alkalinity and also by evaluating simultaneously all kinetic measurements



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(Table S1).62 In agreement with the noted kinetic observations, the results reveal the existence of at least three absorbing species apart from glycine. These are presumably MCG, an intermediate and a final product. The kinetic traces at 228 nm could be fitted to a double exponential function.

A = A1e − k obs1t + A2 e − k obs2 t + A∞

(3)

where kobs1 is the same as in equation 2 and kobs2 characterizes the first-order transformation of the intermediate. The values of kobs1 and kobs2 were typically very similar which introduced an instability problem in fitting some of the experimental traces to equation 3 by using a non-linear least squares fitting routine. Sometimes the algorithm diverged and the rate constants could not be calculated. This problem was circumvented by using the following evaluation protocol. First, the 280 nm trace was fitted to equation 2. The result for kobs1 was substituted into equation 3 and was not allowed to float during the fitting of the traces at 228 nm to equation 3. The calculations yielded well defined kinetic results. The dependence of the pseudo-first-order rate constants on the alkalinity is shown in Figure 3.

Figure 3.

The dependence of the first-order rate constants kobs1 (■) and kobs2 (●) on the

hydroxide ion concentration, I = 1.0 M (NaClO4), T = 25.0 °C.


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In situations when the kinetic observations are consistent with the sum of two exponential terms as in equation 3, there is always a concern regarding the assignment of the two pseudo-first order rate constants. In spite of plausible expectations, it may happen that the smaller rate constant corresponds to the first step in such a reaction scheme.63 This possibility can be eliminated in this system by considering that the only absorbing species is MCG at 280 nm (except, as noted, at high glycine excess). Thus, kobs1 corresponds to the decay of this species and can be assigned to the first step of the overall process. The rate of this step is proportional to the hydroxide ion concentration. The second rate constant is pH independent. The interpretation of kobs1 and kobs2 will be provided together with the mechanistic considerations. The kinetic results are consistent with the report of Hand et al. in that the initial decomposition of MCG is a first order process.45 However, it is also clarified that the noted deviation from the first order behavior in that report was not due to the presence of N,N-dichloroglycine but the formation and further reaction(s) of an intermediate. This effect was overlooked in each of the earlier studies most likely because only the absorbance change at the characteristic MCG band (λmax = 255 nm) was monitored. In this respect, it should be emphasized that only a very narrow wavelength range in the far UV region is suitable for obtaining indisputable evidence for this kinetic pattern. An interesting and highly unusual feature of this system is that while kobs1 and kobs2 are independent from the concentration of glycine, the final absorbance clearly increases by increasing the excess of this reagent. This finding is consistent with a kinetic model which includes two rate determining reaction steps with the formation of intermediates and competing fast reactions in the final stage of the overall process. The relative rates of the competing reaction paths determine the final stoichiometric ratio of the products.



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Mechanistic considerations. The spectral changes observed in the UV region are not suitable for identifying the intermediates and products. Thus, 1H and 13C NMR studies were carried out to obtain detailed information about these species. It should be noted that tuning the NMR instrument for recording time dependent spectra requires several minutes during which the decomposition of MCG progresses and the very beginning of the reaction cannot be monitored by this technique. Typical 1H NMR spectra as a function of time are shown in Figure 4, and the variation of the peak intensities as a function of time for the reactants, an intermediate and products is shown in Figure 5.

Figure 4. Time dependent 1H NMR spectra for the decomposition of MCG in aqueous solution. The left and the right sides of the spectra are shown with different magnifications. CGLY0 = 1.00 × 10-2 M, CMCG0 = 1.00 × 10-2 M, COH− = 0.054 M, T = 25.0 °C.


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Figure 5. The intensities of the 1H NMR peaks as a function of time during the decomposition of MCG. Plots are shown for the peaks of the reactants (a), an intermediate (b) the main product (c) and a minor product (d). The assignment of the peaks corresponds to Figure 4. CGLY0 = 1.00 × 10-2 M, CMCG0 = 1.00 × 10-2 M, COH− = 0.054 M, T = 25.0 °C. The peaks observed at 3.16 and 3.57 ppm are characteristic for the CH2 groups of glycine and MCG, respectively. The time dependencies of the intensities of these peaks are consistent with complete disappearance of MCG and a slight decrease in the concentration of glycine (Figure 5a). The peaks at 3.80 (P1a) and 8.08 (P1b) ppm belong to the main final product of the reaction because they are far bigger than any other new peak in the spectra and their intensities reach a steady value at longer reaction times (Figure 5c). These peaks are consistent with the presence of a CH2 and an aldehyde group, which belong to the same molecule according to the 1H-1H COSY NMR spectrum (Figure S3). The very same peaks were obtained in an aqueous solution of N-formylglycine prepared as described in the literature,57 thus, it was concluded that the main product is most likely this species. The MS spectra with well defined peaks at m/z = 102.022 and 124.006 (when a H+ is replaced by a Na+) in the negative ion mode, and

13

C spectra with

peaks at 48 (CH2), 184 (COO−) and 167 (CHO) ppm corroborate this conclusion (Figure S4).


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These findings are also in accordance with an earlier report on the 1H NMR spectrum of Nformylglycine in D2O with CH2 and CHO peaks at 3.78 – 3.86 and 8.06 ppm, respectively.64 The rotation of the formyl group around the amide bond is hindered in N-formylglycine and this species exists as a mixture of two distinct isomers in aqueous solution as shown in Scheme 1.

Scheme 1. Isomers of N-formylglycine The two smaller peaks in the 1H NMR spectra at 3.84 ppm (P2a, CH2) and 7.91 ppm (P2b, CHO) are consistent with the presence of the second isomer. The time profile for the formation of this compound is shown in Figure 5d. In general, the CH2 and aldehyde proton signals of the cis isomer are expected to show high- and downfield-shifts compared to the trans form, respectively. Thus, the comparison of the chemical shifts of the corresponding peaks reveals that these two smaller peaks belong to the cis isomer.

The ratio of the intensities of the

corresponding peaks are 10:1 (CH2) and 9:1 (CHO) i.e. the trans isomer is present in an about nine-fold excess over the other isomer. In general, the integrals of the 1H NMR peaks around the 4.8 ppm water signal should not be used for quantitative evaluation when a water suppression method is used because a fully selective magnetization of water protons cannot be achieved. Such a problem is not expected with the CH2 peaks of MCG, glycine and N-formylglycine because they are far from the water peak. Thus, it is important to note that the sum of the CH2 intensities of the two isomers of the main product is almost half of the initial CH2 intensity of MCG. This suggests that the formation of one N-formylglycine molecule requires the consumption of two MCG molecules. 15 ACS Paragon Plus Environment


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The NMR spectra feature several other peaks, the identification of which is not straightforward. In order to assign these peaks, MCG and glycine, which is present always in excess, and conceivable intermediates were mixed. In accordance with literature results, we assume that formaldehyde is formed in this system as an intermediate. An earlier study45 and our experimental observations, discussed later in the paper, unequivocally confirm the formation of glyoxylate, too. In independent NMR experiments, we confirmed that the 1H NMR signal of formaldehyde is detectable in water at 3.45 ppm under alkaline conditions (Figure S5). Such a peak was not observed in the 1H NMR spectra of the spent reaction mixtures confirming that formaldehyde does not form as a final product in the decomposition of MCG. In the case of glyoxylate, decent peaks were observed in neutral – slightly alkaline solutions at 5.33 ppm which belong to the hydrated form of glyoxylate in agreement with spectra reported earlier.65, 66 This peak became broader and partly overlapped by the water peak in highly alkaline solution. The peak broadening is most likely due to hydroxide ion catalyzed proton exchange with the solvent. After a few hours of incubation time under alkaline conditions, re-acidification of the sample restored the original peak of glyoxylate confirming that this species does not decompose at high pH. The two small peaks at 4.19 (P3a) and 7.66 (P3b) ppm in the 1H NMR spectra confirm the formation of a minor product. New 1H NMR peaks did not appear in the spectra when glycine and formaldehyde was mixed indicating that the equilibrium between these species and the corresponding imine is probably shifted toward the reactants. However, peaks P3a and P3b immediately appear in the 1H NMR spectrum when glycine is reacted with glyoxylate. Increasing the concentration of glycine increases the intensities of these peaks. When these reactants are mixed in 1:1 ratio, a relatively large glycine CH2 peak still exists in the spectrum


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confirming that a substantial amount of glycine remains intact in the reaction mixture. This observation leads to the conclusion that the reactants are in equilibrium with the product, which is presumably the corresponding Schiff base. In an earlier study, the complex formation of copper(II) with Schiff bases of various aldehydes and ketones was studied in detail and several complexes could be prepared with such ligands.67 The complex of the Schiff base between glyoxylic acid and glycine could not be prepared in that work, but transient existence of this species was indicated. To the best of our knowledge, no characteristic NMR data are available for this Schiff base in the literature except for those of the corresponding diester ethyl (E)-N[(ethoxycarbonyl)methylene] glycinate.68 This compound was reported to exhibit resonances in the 1H NMR spectrum recorded in CDCl3 at 4.10 (CH2) and 7.85 (═CH) ppm and this is in good agreement with our signals observed in H2O. These peaks disappear from the spectra after an extended period of time because the corresponding Schiff base decomposes. On the basis of the combination of the kinetic and NMR results, a detailed mechanism is proposed for the decomposition of MCG. The outline of the mechanism is shown in Scheme 2, while the details of each major segments of the mechanism are shown in Schemes 3 – 5.



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Scheme 2. The outline of the mechanism of MCG decomposition. The two-step kinetic feature shown in Figure 2 challenges the assumption that the decomposition of MCG proceeds via a concerted Grob fragmentation mechanism.45, 47, 49 The first step is first-order with respect to OH− (Figure 3), thus, it is reasonable to assume that the decomposition of MCG is initiated by the formation of a carbanion which loses a chloride ion in a subsequent reaction step. (Scheme 3).

Scheme 3. The initial sequence of the decomposition of MCG (Segment I in Scheme 3). The pKa of the CH2 group of MCG is expected to be high,69

thus the corresponding

equilibrium in Scheme 3 is shifted far to the left. It follows that kobs1 can be expressed as follows: kobs1 = KOH k1[OH−]

(4)


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The fitting of the data in Figure 3 to a straight line yields KOH k1 = (6.4 ± 0.1) ×10-2 M−1s−1 which is in good agreement with results reported in an earlier study made in 0.5 M NaCl (4.4×10−2 M−1s−1).48 Two competing parallel reaction paths can be envisioned for the hydrolysis of the iminoacetate (paths II and III in Scheme 2). Decarboxylation leads to the formation of formaldimine which produces formaldehyde in a subsequent hydrolytic step (Scheme 4). Ammonia was not quantified in this study but the Nessler-test56 confirmed its formation in substantial amount as reported earlier.45, 47-49 Kinetic traces obtained in the decomposition and the reactions of MCG with formaldehyde and glyoxylate are compared in Figure 6. The results at both wavelengths are consistent with fast decay of MCG in the presence of formaldehyde, while the formation of an intermediate was not detected at 228 nm. In accordance with these observations 1H NMR spectra feature only the characteristic peaks of N-formylglycine and glycine which is present in excess. These results lend support to the mechanistic considerations discussed above.

Scheme 4.

The formation of N-formylglycine from the imino-acetate via formaldehyde

(Segment II in Scheme 2).



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Figure 6. The comparison of the kinetic traces for the decomposition of MCG (○) and the reactions of MCG with formaldehyde (◊) and glyoxylate (∆) at 280 (a) and 228 nm (b). CGLY0 = 2.00 × 10-4 M, CMCG0 = 1.00 × 10-3 M, COH− = 0.054 M, CCHOH0 = 2.00 × 10-3 M, CGLYOX0 = 1.00 × 10-3 M, I = 1.0 M (NaClO4), T = 25.0 °C. The other path (Path III, in Scheme 2) for the formation of N-formylglycine is shown in Scheme 5. This scenario takes into consideration that the imino-acetate is in equilibrium with glyoxylate and NH3 via hydroxyglycine.65

As discussed above, equilibrium is established

between glycine and glyoxylate and the corresponding Schiff base is formed. These equilibrium steps are shifted toward the hydrolysis of the imines when glyoxylate is consumed in a subsequent irreversible reaction step with MCG. As shown in Figure 6a, glyoxylate quickly consumes MCG. At 228 nm, the initial sharp increase in absorbance is followed by a slower 20 ACS Paragon Plus Environment

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decay. The kinetic profile is very similar to the one observed in the decomposition of MCG in the absence of added glyoxylate.

The only difference is that the absorbance change is

considerably bigger and the increase is much faster in the glyoxylate – MCG reaction. It follows that this reaction is responsible for the formation of a reactive intermediate in the decomposition of MCG which produces N-formylglycine in further reaction steps.

Scheme 5. The formation of N-formylglycine from the imino-acetate via glyoxylate (Segment III in Scheme 2). 1

H NMR measurements fully support these considerations. When MCG and glyoxylate is

mixed in equivalent concentrations, the CH2 peak of MCG practically disappears by the time when the first spectrum is taken (ca. 6 min). After the consumption of MCG, the intensities of the two peaks assigned to the intermediate steadily decrease while the intensities of the Nformylglycine peaks steadily increase (Figure 7).



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Figure 7.

1

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H NMR spectra recorded after mixing MCG and glyoxylate in equivalent

concentrations.

The left and the right sides of the spectra are shown with different

magnifications. CMCG0 = 2.00 × 10-2 M, CGLYOX0 = 2.00 × 10-2, M CGLY0 = 4.00 × 10-3 M, COH− = 0.054 M, T = 25.0 °C. In a set of related experiments, the reaction was quenched by freezing after about 5 minutes reaction time when the intermediate concentration was expected to reach its maximum. The frozen sample was lyophilized and the 1H NMR spectrum of the solid material was recorded in D2O. This experiment confirmed again, that the reaction of glyoxylate and MCG generates the intermediate which leads to the formation of N-formylglycine. The intensity ratio of the peaks at 4.08 and 9.17 ppm is 2:1. We assign these peaks to the CH2 and NH groups of N-oxalylglycine assuming that the NH proton is immobilized via an internal hydrogen bond with the negatively charged carboxylate group. This assumption is in line with earlier 1H NMR results reporting peaks at 3.8 and 8.97 ppm for N-oxalylglycine in DMSO-d6.70 Evaluation of the kinetic traces (Figure 2b) yields kobs2 = (1.2 ± 0.1) × 10-3 s−1. The simplest explanation would be that this rate constant is characteristic for the hydrolytic decomposition of the imino-acetate. The reaction is irreversible via the formaldehyde path, and if the equilibria are shifted toward the formation of glyoxylate it also can be treated as first order process through the 22 ACS Paragon Plus Environment

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alternative path. However, this explanation is not consistent with the NMR results because does not explain the formation of an intermediate with peaks at 4.08 and 9.17 ppm. Furthermore, the experimental data clearly suggest that the formation of the intermediate requires the presence of glyoxylate, which itself is an intermediate formed from the imino-acetate. Thus, the reaction associated with kobs2 must occur after the reaction step producing glyoxylate. We assume that the intermediate is N-oxalylglycine, which is quickly formed in a reaction sequence shown in Scheme 5 and decomposes relatively slowly in a first-order process characterized by kobs2. In order to confirm this assumption, the absorbance decay in the glyoxylate – MCG reaction at 228 nm (Figure 6b) was fitted to an exponential function using the data only from 500 s. The estimated rate constant, k = (1.2 ± 0.1) × 10−3 s−1, is in excellent agreement with kobs2 obtained from the kinetic studies on the decomposition of MCG.

This lends strong support to the

mechanistic considerations presented here. It is a key issue how the two pathways contribute to the overall reaction. The experimental results provide a straightforward evidence that glyoxylate is an essential intermediate in the decomposition of MCG.

Coherent interpretation of the experimental results would not be

possible without this species.

Consequently, the corresponding reaction path must have a

significant role in the reaction. In contrast, the significance of formaldehyde is questionable even though the formation of this species was postulated as a final product of MCG decomposition in earlier studies. There is a relatively easy way to conceptualize its involvement in the formation of N-formylglycine and if it is formed it reacts with MCG in a relatively fast reaction.

Nevertheless, neither direct nor indirect evidences confirm the formation of this

intermediate. In other words, path II may play a subordinate role only or even can be left out



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from the mechanism without compromising the validity of the interpretation of the experimental results. We consider N-formylglycine as the main product of MCG decomposition. In this respect, it should be emphasized that the reaction time did not exceed a few hours in this work and all considerations apply for this time scale. When the spent reaction mixture was allowed to stand for several days, the main product disappeared from the system and 1H NMR spectra confirmed only the presence of glycinate and formate ions in substantial amounts. In fact, the small formate 1

H NMR peak at 8.43 ppm (P4) showed up in the spectra a few hours after initiating the

decomposition of MCG. This clearly indicated that the main product underwent slow hydrolysis. This reaction does not affect the implications based on the results presented here. CONCLUSIONS This work confirms that the stability of the MCG solution is strongly dependent on the pH. Under alkaline conditions, the rate of the decomposition of MCG is linearly dependent on the hydroxide ion concentration. The decomposition is a rather complex process producing Nformylglycine as the main product. While the conditions used here are different from those in biological systems, the results may be relevant in understanding the biological role of Nchloramines. The decomposition of MCG also proceeds under less alkaline conditions and the “history” of the stock solution may have a strong impact on the experimental observations. The kinetic features of this system are also unique in that the initial rate determining reaction sequence is followed by a series of fast equilibrium steps and the second rate determining step occurs toward the end of the reaction sequence. In contrast to earlier assumptions, the reaction does not proceed according to the concerted Grob mechanism and does not produce formaldehyde as a final product. The main product of


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the reaction is N-formylglycine. There are reports on the biological activities of this species71-76 and the transient formation of N-oxalylglycine, which is known to be an efficient enzyme inhibitor,77-82 can also be the subject of interest. In this respect, it is important to note that Nformylglycine may accumulate in stock solutions of MCG and cause unwanted artefacts in chemical and biological studies. ASSOCIATED CONTENT Supporting Information The second order rate constant for the formation of MCG as a function of pH, spectral changes after mixing glycine with excess hypochlorite, 1H-1H COSY NMR spectrum of the spent reaction mixture in MCG decomposition,

13

C NMR spectrum of the spent reaction mixture in the

decomposition of MCG, 1H NMR spectrum of formaldehyde under alkaline conditions, and singular values of the time-resolved absorbance spectra. This material is available free of charge via the Internet at http://pubs.acs.org. AUTHOR INFORMATION Corresponding Author *

Phone: +36 52 512900 ext.: 22378, Fax: +36 52 518-660, E-mail: [email protected]

Funding This research was funded by the Hungarian Science Foundation under grant no NK 105156, and co-financed by the European Social Fund under the project ENVIKUT (TAMOP-4.2.2.A11/1/KONV-2012-0043) Notes, The authors declare no competing financial interest. Acknowledgement



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The assistance of Mr. Tibor Nagy in the MS experiments is highly appreciated. ABBREVIATIONS MCG, N-chloroglycine; NMR, nuclear magnetic resonance spectroscopy; kobs1, kobs2, pseudo-first-order rate connstants of the observed absorbance change; A, A∞ time dependent and final absorbance of the kinetic traces; A0, A1, A2, amplitudes of the absorbance change in the kinetic experiments.

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Kinetics and mechanism of degradation of a cyclic hexapeptide (somatostatin analogue) in aqueous solution.

Radiolytic decomposition of ciprofloxacin using γ irradiation in aqueous solution.

Factors influencing decomposition rate of amitriptyline hydrochloride in aqueous solution.

Stability and kinetics of degradation of imipenem in aqueous solution.

Hydrolysis and epimerization kinetics of hetacillin in aqueous solution.

Decomposition of amitriptyline hydrochloride in aqueous solution: identification of decomposition products.

Photochemical kinetics of pyruvic acid in aqueous solution.

Kinetics and mechanism of (•)OH mediated degradation of dimethyl phthalate in aqueous solution: experimental and theoretical studies.

Decomposition and Mineralization of Dimethyl Phthalate in an Aqueous Solution by Wet Oxidation.

Speciation and structure of tin(II) in hyper-alkaline aqueous solution.

Mechanism of cellulose gelation from aqueous alkali-urea solution.

Speciation and the structure of lead(II) in hyper-alkaline aqueous solution.

Kinetics and mechanisms of depolymerization of alginate and chitosan in aqueous solution.

Comparative stability of cephalosporins in aqueous solution: kinetics and mechanisms of degradation.

Ab initio kinetics and thermal decomposition mechanism of mononitrobiuret and 1,5-dinitrobiuret.

Thermal degradation kinetics and decomposition mechanism of PBSu nanocomposites with silica-nanotubes and strontium hydroxyapatite nanorods.

Research on the Thermal Decomposition Reaction Kinetics and Mechanism of Pyridinol-Blocked Isophorone Diisocyanate.

The oxidation of tiron by superoxide anion. Kinetics of the reaction in aqueous solution in chloroplasts.

Misoprostol dehydration kinetics in aqueous solution in the presence of hydroxypropyl methylcellulose.

Kinetics of the disordered chain-to-beta transformation of poly(L-tyrosine) in aqueous solution.

Combined capillary electrophoresis and high performance liquid chromatography studies on the kinetics and mechanism of the hydrogen peroxide-thiocyanate reaction in a weakly alkaline solution.

Adsorption of Phenol from Aqueous Solution Using Lantana camara, Forest Waste: Kinetics, Isotherm, and Thermodynamic Studies.

Kinetics and thermodynamic studies for removal of acid blue 129 from aqueous solution by almond shell.

Equilibrium and kinetics of phosphorous adsorption onto bone charcoal from aqueous solution.