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Official Journal of the Japan Wood Research Society

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Stereo-preference in the degradation of the erythro and threo isomers of β-O-4-type lignin model compounds in oxidation processes III: in the reaction with chlorine- and manganese-based oxidants

Abstract

We examined which isomer, the erythro or threo, of a non-phenolic β-O-4-type lignin model compound is stereo-preferentially oxidized in hypochlorite, chlorite, manganese dioxide, or permanganate systems. No clear stereo-preference was observed in the hypochlorite or chlorite systems at alkaline, neutral, or acidic pH except for a slight threo stereo-preference that appeared after the step-wise addition of the oxidizing reagent had been terminated in the neutral or acidic chlorite systems. A clear threo stereo-preference was observed in the manganese dioxide or permanganate systems.

Introduction

Because the β-O-4-type is the most abundant substructure in lignin, cleavage of the β-O-4 bond always controls the delignification and depolymerization of lignin in chemical processes. The diastereomeric erythro (E) and threo (T) isomers exist in the side chain of the β-O-4-type substructure (Fig. 1). These isomers show different reactivities in various chemical reactions. An example reaction is the β-O-4 bond cleavage under alkaline pulping conditions, where the β-O-4 bond of the E isomer cleaves more rapidly than that of the T isomer [1,2,3,4,5,6]. However, only a few papers have examined stereo-preferential degradation in oxidation processes [7,8,9,10]. Some of them reported stereo-preferential degradation of the T isomer [7, 8], while no clear stereo-preference was observed in the others [9, 10].

Fig. 1
figure1

Structures of the β-O-4-type lignin subunit, compound 1E, and compound 1T

Our previous reports examined the stereo-preferential degradation of the E or T isomer, when each isomer of non-phenolic β-O-4-type lignin model compounds was individually oxidized by various radical species generated as active oxygen species under oxygen bleaching conditions [11] or by hydroxyl radical and/or its conjugate base, oxyl anion radical, under alkaline hydrogen peroxide bleaching conditions [12]. A slight stereo-preferential degradation of the T isomer was observed at a high pH (> 13). When the electric repulsion between the side chains of the lignin model compounds and a negatively charged active oxygen species, oxyl anion radical, working at high pH, was quenched by lowering the pH (< 13) or structurally modifying the lignin model compounds, a slight stereo-preferential degradation of the E isomer was observed. We concluded, therefore, that the essential stereo-preference of oxyl anion radical is slightly in favor of the E isomer although the reverse stereo-preference can appear depending on the reaction conditions.

In this paper, the stereo-preferential degradation of the E or T isomer was examined when each isomer of the most general non-phenolic β-O-4-type lignin model compound, 2-(2-methoxyphenoxy)-1-(3,4-dimethoxyphenyl)propane-1,3-diol (1E or 1T, respectively, Fig. 1), was oxidized in hypochlorite, chlorite, manganese dioxide, or permanganate systems.

Materials and methods

Materials

All chemicals except compounds 1E and 1T were purchased from Wako Pure Chemical Industries, Ltd. (Osaka, Japan), Tokyo Chemical Industry Co., Ltd. (Tokyo, Japan), and Sigma-Aldrich Co. LLC. (St. Louis, MO, USA), and used without further purification. Ultrapure water (Puric-Z, Organo Co., Tokyo, Japan) was used in all the experiments.

The synthesis of compounds 1E and 1T and their separation were described in our previous reports [11,12,13]. The structures and purities were confirmed by nuclear magnetic resonance spectroscopy (1H-NMR and 13C-NMR, JNM-A500, JEOL Ltd., Tokyo, Japan). The spectral peaks were indicated in our previous report [11].

Oxidation in hypochlorite systems

Either compound 1E or 1T (30 µmol) or both (15 µmol each) were dissolved in 30 mL of water at an alkaline pH of 13.3, neutral pH of 6.3, or acidic pH of 1.0, adjusted by sodium hydroxide, no reagent, or sulfuric acid (H2SO4), respectively, in a round-bottom glass flask (50 mL volume). The flask was placed in a water bath and heated to 70 °C in the alkaline or neutral reaction, while the acidic reaction was conducted at room temperature. Sodium hypochlorite (NaClO, 0.30 mmol, 10 times mole amount of compound 1E or 1T) was added to the reaction solution to initiate the reaction. The same amount of NaClO was added a total of five times at constant intervals of 10 min, so that the total amount of NaClO added was 1.5 mmol. The reactions were terminated at a reaction time of 60 min (20 min after the final NaClO addition).

A portion of the reaction solution was withdrawn at prescribed reaction times to quantify the surviving compound 1E or 1T and to measure the pH. The detailed workup procedures for quantification were the same as those described in our previous reports [11,12,13] except for neutralizing the acidic reaction solution with sodium hydrogen carbonate.

Oxidation in chlorite systems

The reaction was as described above for the NaClO oxidation except that: (1) sodium chlorite (NaClO2) was employed instead of NaClO, maintaining the molar ratio to compound 1E or 1T, and (2) the temperature was 70 °C for all reactions.

Oxidation in a manganese dioxide system

Commercially available manganese dioxide (MnO2, Wako Pure Chemical Industries, Ltd.) was ground into powder on a mortar and iodometrically titrated, revealing that the oxidation power of the MnO2 powder was 89% of the theoretical value. The MnO2 powder (26 mg, 0.30 mmol) was aged in 25 mL of sulfate buffer solution (0.50 mol/L, pH 1.0) for 120 min in a round-bottom glass flask (50 mL) at room temperature. Either compound 1E or 1T (6.0 µmol) was dissolved in 5.0 mL of another sulfate buffer solution (0.50 mol/L, pH 1.0), and this buffer solution was added to the buffer solution containing the aged MnO2 powder to initiate the reaction. The reaction was conducted at room temperature for 350 min. The initial concentration of compound 1E or 1T was 0.20 mmol/L, giving a molar ratio to the MnO2 powder of 1/50.

A portion of the reaction solution was withdrawn to quantify the surviving compound 1E or 1T at prescribed reaction times. The withdrawn solution was immediately extracted with ethyl acetate (EtOAc) containing an internal standard compound, 3,4,5-trimethoxybenzaldehyde. The water layer was further immediately extracted twice with EtOAc. The combined EtOAc layer was washed with a saturated sodium hydrogen carbonate solution followed by brine, and concentrated under reduced pressure. An aqueous 50% methanol (CH3OH) solution (v/v) was added to the residue and filtrated with a membrane filter. The obtained mixture was analyzed by a high-performance liquid chromatograph (HPLC, LC-10A, Shimadzu Co., Kyoto, Japan) equipped with an SPD-M10A detector (Shimadzu Co.), using the absorbance at 280 nm.

Conditions of HPLC were as follows. Column: Luna 5u C18 (2) 100A (150 mm × 4.6 mm, Phenomenex Inc., Torrance, CA, USA); oven temperature: 40 °C; flow rate: 0.2 mL/min; solvent system: gradient CH3OH/H2O (v/v) from 30/70 to 40/60 for 30 min and maintained for 15 min, total time 45 min.

Oxidation in permanganate systems

As an acidic reaction, either compound 1E or 1T was dissolved in 24.6 mL of sulfate buffer solution (0.50 mol/L, pH 2.0). To this buffer solution was added 0.40 mL of 20 mmol/L potassium permanganate (KMnO4) solution to initiate the reaction. The reaction was conducted at room temperature. The initial concentrations of compound 1E (or 1T) and KMnO4 were 0.20 and 0.32 mmol/L, respectively.

As a neutral reaction, either compound 1E or 1T was dissolved in 23 mL of water (pH 6.3). To this solution was added 2.0 mL of 20 mmol/L KMnO4 solution to initiate the reaction. The reaction was conducted at room temperature. The initial concentrations of compound 1E (or 1T) and KMnO4 were 0.20 and 1.6 mmol/L, respectively.

A portion of the reaction solution was withdrawn to quantify the surviving compound 1E or 1T at prescribed reaction times. The detailed workup procedures for quantification and conditions for the HPLC analysis were the same as described above for the MnO2 system.

Results and discussion

Oxidation in hypochlorite systems

Figure 2a shows the degradation of compound 1E or 1T when either compound was individually (twice each) or both compounds were together (once) reacted at an alkaline pH of 13.3 and 70 °C. Any observed degradation was always less than 5% at the end of the reaction (60 min). No clear difference was observed in the degradations between compounds 1E and 1T. An iodometric titration confirmed almost no consumption of the oxidation power at the end of the reaction. Because hypochlorite anion (ClO¯) rather than hypochlorous acid (HClO) existed as the major species due to the pKa value of HClO (7.58 at 20 °C [14]), the result indicates that ClO¯ is rather stable under the conditions and compounds 1E and 1T do not have any good reaction sites with ClO¯.

Fig. 2
figure2

Time courses of the changes in the recovery yields of compounds 1E and 1T, when either compound was individually or both compounds were together treated in the hypochlorite system at a pH of: a 13.3 and 70 °C or b 6.3 and 70 °C. Compound 1E: filled circle, filled diamond, filled star; compound 1T: open circle, open diamond, open star; in the individual reaction of compound 1E or 1T: filled circle, filled diamond, open circle, open diamond; in the reaction of both compounds 1E and 1T together: filled star, open star

Figure 2b shows the degradation of compound 1E or 1T when either compound was individually (once each) or both compounds were together (once) reacted at a neutral pH of 6.3 and 70 °C. Any observed degradation was rapid, with complete disappearance at a reaction time of about 40 min. No clear difference was observed in the degradations between compounds 1E and 1T. HClO rather than ClO¯ existed as the major species due to the pKa value. A possible reaction mechanism is nucleophilic attack of a nucleophile on the Cl atom of HClO, which is essentially the reaction of chlorine cation (Cl+) with the nucleophile. The aromatic ring of compound 1E or 1T is a more plausible candidate for the nucleophile than the side chain portion. This would explain why no clear stereo-preference was observed in the reaction with compound 1E or 1T, each of which has a stereo-structurally different side chain. Both compounds 1E and 1T were stable without the addition of NaClO under these conditions.

Both compounds 1E and 1T were together reacted at an acidic pH of 1.0 and room temperature (twice). Because the degradation was quite rapid in a preliminary experiment, room temperature was employed (data not shown). The degradation was still fast, and the recovery yields were 6 and 3%, respectively, at the first sampling time (10 min). This result indicates that compounds 1E and 1T were efficiently oxidized. However, the difference was not large enough to safely state that compound 1T was stereo-preferentially degraded rather than compound 1E. Generally, molecular chlorine (Cl2) is generated when HClO reacts with hydrogen chloride (HCl) at a low pH. Because H2SO4 was used for the pH adjustment, little generation of Cl2 took place. A nucleophilic attack of the aromatic ring to the Cl atom of HClO would be a possible major reaction, similar to the above-described reaction at the neutral pH. However, this attack would be more rapid than that at the neutral pH due to the protonation of HClO at the acidic pH (H2O+Cl) and consequent increase of the electrophilicity of the Cl atom. This would result in the efficient oxidation of compounds 1E and 1T and a lack of stereo-preferential degradation. Both compounds 1E and 1T were rather stable without the addition of NaClO under these conditions.

Oxidation in chlorite systems

Figure 3a shows the degradation of compound 1E or 1T when either compound was individually (once each) or both compounds were together (once) reacted at an alkaline pH of 13.3 and 70 °C. Any observed degradation was less than 5% at the end of the reaction. No clear difference was observed in the degradations between compounds 1E and 1T. An iodometric titration confirmed almost no consumption of the oxidation power at the end of the reaction. Because chlorite anion (ClO2¯) rather than chlorous acid (HClO2) existed as the major species due to the pKa value (1.94 at 25 °C [15]), the result indicates that ClO2¯ is rather stable under these conditions and compounds 1E and 1T do not have any good reaction sites with ClO2¯.

Fig. 3
figure3

Time courses of the changes in the recovery yields of compounds 1E and 1T, when either compound was individually or both compounds were together treated in the chlorite system at a pH of: a 13.3 and 70 °C, b 6.3 and 70 °C, or c 1.0 and 70 °C. Compound 1E: filled circle, filled diamond, filled square, filled star; compound 1T: open circle, open diamond, open square, open star; in the individual reaction of compound 1E or 1T: filled circle, filled diamond, filled square, open circle, open diamond, open square; in the reaction of both compounds 1E and 1T together: filled star, open star

Figure 3b shows the degradation of compound 1E or 1T when either compound was individually reacted at a neutral pH of 6.3 and 70 °C (thrice each). The degradation was relatively rapid until a reaction time of about 40 min and accelerated during this period. (This period is hereafter described as ‘the period of rapid degradation’.) After the final addition of NaClO2 at a reaction time of 40 min, slower but continuous degradation was observed until the reaction was terminated at a reaction time of 60 min (‘the period of slower degradation’). No clear difference was observed in the degradations between compounds 1E and 1T in the period of rapid degradation, while a difference gradually appeared with slight stereo-preferential degradation of compound 1T in the period of slower degradation. It is generally considered that chlorine dioxide (ClO2) is generated by the disproportionation of ClO2¯ or other similar reactions as the most active species, although other chlorine-related oxidants are also produced. On the basis of the obtained results, we suggest that ClO2 as well as other chlorine-related oxidants attacks compound 1E or 1T without any clear stereo-preference in the period of rapid degradation and the formation of ClO2 as well as other chlorine-related oxidants from ClO2¯ (rather than from HClO2 due to the pKa value) accelerates with the progress of the reaction. Because the major oxidant ClO2 is a radical and preferably adds to aromatic nucleus of an aromatic substrate including lignin model compounds as an electrophile accompanied by the liberation of ClO2¯ or HClO2 (depending on pH) and generation of an aromatic radical cation [16,17,18,19], it is natural that no clear stereo-preference was observed in the period of rapid degradation. It is unclear, on the other hand, why the slight T stereo-preference was observed in the period of slower degradation. Some oxidants generated from ClO2¯, ClO2, and others would still exist in the period of slower degradation, and show the slight T stereo-preference. Both compounds 1E and 1T were stable without the addition of NaClO2 under these conditions.

Figure 3c shows the degradation of compound 1E or 1T when either compound was individually reacted at an acidic pH of 1.0 and 70 °C (thrice each). Although the degradation of each compound was similar to that observed at the neutral pH, the following observations were different from those at the neutral pH. (1) The degradation was slightly slower than that at the neutral pH, and acceleration of the degradation was not observed in the period of rapid degradation. (2) The degradation of compound 1T was slightly greater than that at the neutral pH in the period of slower degradation, and the difference in the degradations between compounds 1E and 1T was larger than that at the neutral pH. Observation (1) may indicate that the generation of ClO2 does not exactly follow the same mechanism as that at the neutral pH. HClO2 was the main species rather than ClO2¯ at the acidic pH due to the pKa value, which might result in different formation mechanisms of ClO2 from that at the neutral pH. Observation (2) suggests that the profile of oxidants responsible for the degradation of compound 1E or 1T was not exactly the same as that at the neutral pH. In accordance with the expectation on the basis of the description in the previous paragraph, no clear stereo-preference was observed in the period of rapid degradation. Both compounds 1E and 1T were rather stable without the addition of NaClO2 under these conditions.

Oxidation in a manganese dioxide system

Figure 4 shows the degradation of compound 1E or 1T when either compound was individually reacted at a pH of 1.0 and room temperature (thrice each). The degradation of compound 1T was clearly more rapid than that of compound 1E, which indicates that the stereo-preference of MnO2 is the T isomer. Because MnO2 oxidized compound 1E or 1T as large precipitating aggregates, it can be understood that the stereo-structure of the side chain affects the reaction rate. It is presumed that MnO2 aggregates are able to approach compound 1T more easily than compound 1E due to the stereo-structure of the side chain. The obtained results, however, are not sufficient to rule out the possibility that the stereo-preference of MnO2 observed here is dependent on a particular structure of the aggregates and that other MnO2 aggregates prepared by other methods show the reverse or almost no stereo-preference.

Fig. 4
figure4

Time course of the change in the recovery yield of compound 1E or 1T, when either compound was individually treated in the manganese dioxide system at a pH of 1.0 and room temperature. Compound 1E: filled circle, filled diamond, filled square; compound 1T: open circle, open diamond, open square

Oxidation in permanganate systems

Figure 5a, b shows the degradation of compound 1E or 1T when either compound was individually reacted at a pH of 2.0 or 6.3, respectively, and room temperature. Because the degradation was too rapid to follow at a pH of 1.0 using a large excess of KMnO4 in a preliminary experiment, the applied pH was 2.0 and the molar ratio of KMnO4 to compound 1E or 1T was 1.6 in the acidic system. The molar ratio of KMnO4 to compound 1E or 1T was 8.0 at a pH of 6.3.

Fig. 5
figure5

Time course of the change in the recovery yield of compound 1E or 1T, when either compound was individually treated in the permanganate system at a pH of: a 2.0 and room temperature or b 6.3 and room temperature. Compound 1E: filled circle; compound 1T: open square

The degradation of compound 1T was clearly greater than that of compound 1E at a pH of 2.0. It is possible to assume that MnO2 was generated as an intermediate at this pH and co-oxidized compound 1E or 1T, although the half-reaction of permanganate (MnO4¯) in an acidic medium is described as: MnO4¯ + 8H+ + 5e¯ → Mn2+ + 4H2O. Because the MnO2 oxidation of compound 1E or 1T was confirmed to be much slower than the observed oxidation rate at the pH, the observed T stereo-preference is concluded to be due to the reaction of MnO4¯ at the pH without the effect of MnO2. Because the oxidation system was homogeneous differently from that of MnO2 and it is commonly believed that the oxidation mechanism of MnO4¯ is the attack on an electron rich double bond, aromatic nucleus, etc., the effect of the stereo-structure of the side chain on the MnO4¯ oxidation cannot easily be explained. The large size of MnO4¯ may possibly relate to the T stereo-preference, resulting in the side chain of compound 1T being more ready than that of compound 1E to interact with MnO4¯.

The degradation of compound 1T was also greater than that of compound 1E at a pH of 6.3, although the degradation was not great. The half-reaction of MnO4¯ in a neutral medium is described as: MnO4¯ + 4H+ + 3e¯ → MnO2 + 2H2O, which explains the observed slow oxidation at this pH due to the requirement for H+. Because the degradation of compound 1E or 1T was not great and was observed only at an early stage of the reaction in spite of the application of 8 times mole amount of KMnO4, some degradation products might have been oxidized by MnO4¯ more easily than compound 1E or 1T. The T stereo-preference is not easily explainable, either.

Conclusions

The erythro (E) or threo (T) isomer of the most common non-phenolic β-O-4-type lignin model compound was individually or both isomers together oxidized in hypochlorite, chlorite, manganese dioxide, or permanganate systems to examine which isomer was stereo-preferentially oxidized. No stereo-preference was observed in any hypochlorite or chlorite system at any applied pH except in the later stage of the chlorite system at neutral or acidic pH, where the stereo-preferential degradation of the T isomer was observed in this later stage. Stereo-preferential degradation of the T isomer was also observed in the manganese dioxide and permanganate systems. These stereo-preferences cannot easily be explained, although some of them may possibly result from the ease of approach of the oxidant to the T isomer rather than to the E isomer.

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Correspondence to Tomoya Yokoyama.

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Posoknistakul, P., Akiho, S., Akiyama, T. et al. Stereo-preference in the degradation of the erythro and threo isomers of β-O-4-type lignin model compounds in oxidation processes III: in the reaction with chlorine- and manganese-based oxidants. J Wood Sci 64, 451–457 (2018). https://doi.org/10.1007/s10086-018-1714-z

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Keywords

  • Chlorine dioxide
  • Chlorite
  • Hypochlorite
  • Manganese dioxide
  • Permanganate