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Phase-Change Absorption of SO2 by Imidazole in Organic Solvents and Conversion of the Absorption Product in the Presence of Water and Oxygen Yang Wang, Wenbo Zhao, Muyuan Chai, Genming Li, Qingming Jia, and Yuan Chen Energy Fuels, Just Accepted Manuscript • DOI: 10.1021/acs.energyfuels.7b02694 • Publication Date (Web): 20 Nov 2017 Downloaded from http://pubs.acs.org on November 21, 2017

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Phase-Change Absorption of SO2 by Imidazole in Organic Solvents and Conversion of the Absorption Product in the Presence of Water and Oxygen Yang Wang, Wenbo Zhao*, Muyuan Chai, Genming Li, Qingming Jia, Yuan Chen Faculty of Chemical Engineering, Kunming University of Science and Technology, Kunming, 650500, China E-mail: [email protected] Abstract: Conventional removal of SO2 yields a large amount of waste salts or requires significant energy for regeneration of the absorbent. Phase-change capture has been considered a possible method to solve these problems since only the absorption product needs to be disposed for recovery and the solvent could be reused directly. In the present work, the phase-change absorption behavior of 1,3,5-trimethylpyrazole, 2-methylimidazole, 1-methylimidazole, and 1,2-dimethylimidazole has been investigated in organic solvent. Among these absorbents, only 1,2-dimethylimidazole exhibited an obvious phase-change performance in solvents three glycol dimethyl ether, propylene carbonate and dimethylacetamide, but not in butanol. The composition and structure of the absorption product were not related to the type of solvent but affected by the air. 1,2-dimethylimidazole reacted with SO2 to form a charge-transfer complex, which converted to pyrosulfite (C5H9N2)2S2O5 in the presence of water. The absorption product was further oxidized to dithionate (C5H9N2)2S2O6 by oxygen from the air. Once the filtrate, which was created from the separation process of the phase-change product and solvent, was exposed to air for several days, another type of oxidative product (C5H9N2)2SO4 was discovered. 1. Introduction The ever-increasing demand of energy results in the extensive burning of fossil fuels, which causes the excessive emissions of SO2. The emissions not only cause adverse impacts on human health but are also corrosive to soil and buildings. Flue gas desulfurization (FGD) is considered as an effective solution to reduce the emissions of SO21. Among many commercial processes, limestone scrubbing is the most widely adopted techniques for SO2 removal in the energy industry2. However, this method has some inherent drawbacks, including large amounts of wastewater and waste salts (calcium sulfate), complicated technology, and fine particle pollution. Reusable absorbents aqueous amines are also commonly used for SO2 removal3; however, their high water content consumes a large amount of energy for regeneration of the absorbent. Therefore, novel solvents or liquid materials for SO2 capture are extremely desired. Ionic liquids (ILs) are considered as a potential candidate due to unique properties such as low vapor pressure, high thermal stability and tunable structure. Up to now, many ILs have been investigated for the capture of SO2 such as guanidinium-based4, imidazolium-based5, pyridinium-based6, phosphonium-based7, piperazinium-based8 and hydroxyl ammonium based ILs9. On the other hand, carboxylate anions10, aprotic heterocyclic anions11, nitrile anions, thiocyanate anions12-14 and phenolate anions15 were also found to have the ability to capture SO2. Wang et.al have reported that the azole-based ILs with multiple active sites on anions exhibited an

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extremely high absorption capacity of SO2 (>3.5 mol/mol)16. Their SO2 capture ability could be further improved by grafting a tertiary amino group onto the cation17. Furthermore, ether-functionalized ILs also exhibit high capacity of SO2 capture and excellent reversibility through a combination of chemical and physical absorption from the tetrazolate anion and ether functional group18. Recently, the alcohol hydroxyl group was also proposed as a means to capture SO2 in the presence of a strong base such as tertiary amine and guanidine19. The combination of alcohol and base into one molecule would reduce the volatility of the bimolecular system20. Generally, chemical absorptions with high capacity always release high absorption enthalpy, which results in high energy requirements for regeneration. The method to decrease the basicity of ILs by tuning substituent of anion has been proposed to reduce the interaction between anions and SO2, and hence reduce the energy demand for regeneration. However, there was often a corresponding decrease in absorption capacity when using this method. Therefore, the development of an alternative method to reduce the regeneration energy demand is of particular interest. In the research field of CO2 capture, phase-change absorption as a possible strategy has recently been proposed. In this system, the homogeneous solution was turned into a CO2-lean phase and a CO2-rich phase upon simple bubbling with CO221, 22. The CO2-lean phase can be used directly. Consequently, the energy demand was significantly reduced since only the CO2-rich phase requires heating for regeneration. Usually, a phase-change absorbent consisted of two components, one as the absorbent could capture CO2 and separate from the system as CO2-rich phase, and the other as solvent have the function to adjust the viscosity and mass transfer property of the system. At present, the reported absorbents included amino acid, lipophilic amine, 1,3-dipropil-methyl-xanthine, 3-(methylamino)propylamine, alkanolamine, polyamine, and the related solvents included water, alcohol, ILs23-26. Recently, liquid-solid phase-change behavior of diethylenetriamine for CO2 capture has been reported by our team. The precipitate from CO2 capture can easily be separated from the liquid counterpart27. Unlike the phase-change absorption of CO2, the absorption of SO2 using a similar method has rarely been reported in the literatures. In our previous work, we investigated the phase-change absorption of SO2 by polyamine. The experimental results indicated that diamines such as ethanediamine and piperazine would precipitate from the absorption solution by reacting with SO2 and water from the air to form sulfite crystals28. But hydrophobic triethylene diamine would absorb two SO2 molecules to form a charge-transfer complex even in the presence of water29. In Shannon’s work30, a variety of N-functionalized imidazoles were investigated for absorption of SO2 since they can be readily synthesized at reasonably large scale with simple chemical reactions using inexpensive starting materials31. It was found that some imidazoles were converted from a liquid to a solid or gel upon capturing SO230. The absorption enthalpy for physical dissolution (−4 to −13 kJ/mol) was determined via thermodynamic relationships and the complexation binding energies (−35 to −42 kJ/mol) were determined via density functional theory calculations. However, the composition of the absorption product was not identified by characterization, but was simply deduced to be an SO2•imidazole complex by quantum chemistry simulation. In the present work, some organic solvents such as three glycol dimethyl ether, propylene carbonate, dimethylacetamide and butanol were introduced into the absorption system to adjust the volatility and viscosity of the absorbent imidazoles. The capture product was also found out based on the in-situ FTIR, XRD, and elemental analysis results. Furthermore, the conversion of the absorption

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product in the presence of water and oxygen was also clarified. 2. Experimental 2.1 Materials All chemicals in the present work were purchased from commercially available sources and used as received without further purification. Analytical reagents 1,3,5-trimethylpyrazole (97%), 2-methylimidazole (98%), 1-methylimidazole (99%), 1,2-dimethylimidazole (DMI) (98%), propylene carbonate (99%), dimethylacetamide (99%), and three glycol dimethyl ether (99%) were purchased from Aladdin-Reagent Company in Shanghai. Butanol (99%) was obtained from Tianjin Kemiou Chemical Reagent Co. Ltd. and SO2 (99.9%) was purchased from Chengdu Hongjin Chemical Co., Ltd.

2.2 Characterization FTIR spectra were recorded on a Bruker TENSOR27 Fourier-transform infrared spectrometer. In order to avoid the influence of air, the spectroscopic study for the absorption product was carried out in an in-situ pool with the protection of N2 gas. Powder XRD of the solid was measured using a Rigaku D/max-A X-ray diffractometer with λ= 0.1541 nm, Cu Ka radiation in the 2θ range of 10~90o at room temperature. The CHNS contents of the product were determined using an elemental analyzer (PerkinElmer 2400). TG and DSC were measured on a NETZSCH STA-449 F3 under a flow of nitrogen at a 5 K/min heating rate up to 823 K. Single crystal X-ray diffraction data were collected on a Bruke Apex II CCD diffractometer equipped with CCD area detector using graphite monochromated Mo Ka radiation (λ= 0.71073 Å). The structures were solved by direct methods and refined by a full-matrix least squares analysis using anisotropic thermal parameters for non-hydrogen atoms with the SHELXTL 97 program. All hydrogen atoms were calculated at idealized positions and refined with the riding models. 2.3 Absorption of SO2 The SO2 absorption experiment was carried out in a 250 mL three-necked flask at room temperature. In a typical process, 3 g absorbent and 7 g solvent were put into the flask and then pure SO2 gas, which was first saturated by the corresponding solvent, was bubbled into the reactor at the flow rate of 100mL/min with continued magnetic stirring. After the precipitate formed, the mixture of solid and liquid was separated by filtration. The solid was washed with solvent and dried in a vacuum oven to eliminate the influence of organic solvents. 2.4 Preparation of the single crystal A single crystal suitable for the X-ray diffraction analysis was prepared by dissolving the solid powder in anhydrous methanol and then evaporating the solution in a vacuum oven for a certain time at room temperature. The other single crystal was obtained by evaporating the filter liquor under air for 7 days. Crystallographic data have been deposited in the Cambridge Crystallographic Data Center with CCDC reference numbers 1569803 and 1569812. 3. Results and discussion 3.1 Selection of the absorbent and solvent To discover the optimal system for SO2 absorption, different combinations of absorbents and organic solvents were investigated as shown in Table 1. There was no phase-change phenomenon observed with 1,3,5-trimethylpyrazole as the absorbent no matter which solvent was used. The possible reason may be that its basicity is much lower than that of imidazole due to the

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electron-withdrawing inductive effects of the adjacent N atom. 2-methylimidazole is a white solid at room temperature. It could not be dissolved into propylene carbonate and three glycol dimethyl ether with the mass content of 30%. In the solvents butanol and dimethylacetamide, 2-methylimidazole could be dissolved completely, but a phase-change phenomenon again was not observed. Using 1-methylimidazole as the absorbent, no white solid was formed in four kinds of solvents except three glycol dimethyl ether. On the other hand, using DMI as the absorbent, regardless of the solvent used, a white solid always formed after SO2 was bubbled into the solution. The possible reason may be that its basicity was higher than others, which comes from the electron-donating effect of two methyl. It strengths the interactions between functionalized imidazole and SO2. Some works has shown that electron-donating groups reinforced imidazole interactions with H+ and metal cations32. It should be noted that the amount of precipitate formed in butanol was much lower than that in other solvents. The aforementioned experiments indicated that the occurrence of phase-change phenomenon not only related to the absorbent, but also related to the solvent. On the basis of our work, DMI is a potential candidate for the phase-change capture of SO2. Therefore, only the product from DMI absorbent was investigated in the following work. 3.2 Characterization In order to determine the composition of the absorption product, powder XRD characterization was carried out at first, and the results are shown in Figure 1. The precipitate from butanol was too small to perform XRD analysis and then there is no its XRD pattern. The three similar XRD patterns indicated that solvents have no effect on the composition of the product. In the experimental process, we found that the product exhibited drastic hydroscopicity. Once the absorption product was exposed to the air for several minutes, its XRD pattern was observed to be different from the original one, as shown in Figure 2. Some bands such as at 11.5°, 12.8°, 16.7°, 20.5°, 21.9° and 24.6° disappeared. This phenomenon suggested that the absorption product could absorb water from the air to form another kind of compound. Considering this conversion process, some water was added into the absorption solution to obtain a relatively stable phase-change absorption product. The molar content of water in the solution was equal to that of DMI. Figure 3 shows the XRD patterns of the absorption product from different solvents in the presence of water. The result also indicated that the composition of the product was not affected by solvents since the spectra of product from different solvents have similar patterns. However, their compositions cannot be identified by comparing with the standard spectra in the Jade database. In order to ascertain the composition of the absorption product with or without water, in-situ FTIR was carried out. Figure 4A shows the spectra of DMI. After 2% SO2 gas was introduced into the in-situ pool, the characteristic bands of gas state SO2 appeared at 1161 and 1357 cm-1 (See Figure 4B). In addition, some other new bands appeared at the same time. Once the SO2 gas was replaced by purge gas N2, the characteristic bands belonging to gas state SO2 disappeared, but the bands at 954, 1031, 1051, 1100, 1208 and 1614 cm-1 still existed (See Figure 4C). Among them, the bands at 1100 and 1208 cm-1 can be attributed to the symmetric and asymmetric stretching vibration of SO2 in the organic complex33, 34. This result indicated that SO2 reacted with DMI to form a charge-transfer complex as shown in chemical equation 1. C5H8N2 + SO2 → C5H8N2•SO2 (1) Although there are two N atoms in the DMI molecule, one N atom participated in the conjugation of heterocycle and then only the other one could coordinate with one SO2 molecule. A similar 4


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reaction between amine and SO2 to form charge-transfer complex has been reported in the previous literatures35. After some moisture was introduced into the in-situ pool by N2 gas, the FTIR spectrum of charge-transfer complex changed quickly and significantly. The characteristic bands of charge-transfer complex at 1100 and 1208 cm-1 disappeared and some new bands appeared at 1196, 1174, 1096, 1058 and 969 cm-1 (See Figure 4D). They are the characteristic bands of the S2O52anion and belong to two types of S-O symmetric stretching vibrations36. This result suggested that the charge-transfer complex had converted to pyrosulfite. The possible reaction equation for the charge-transfer complex and water is as follows: 2C5H8N2•SO2 + H2O → (C5H9N2)2S2O5 (2) This deduction was further supported by elemental analysis results. As shown in Table 2, the measured contents of C, H, N and S in the reaction product of DMI and SO2 in the presence of water was 35.76, 5.32, 16.84, and 18.17%, which was consistent with the theoretical values (35.49, 5.36, 16.56, and 18.85%). It should be noted that this product was different from the reaction product of diamines with SO2 in the presence of water, which is sulfite or charge-transfer complex28. To the best of our knowledge, this is the first observation that an SO2 charge-transfer complex reacted with water to form pyrosulfite. We also attempted to ascertain the composition of the product from the reaction of DMI and SO2 without water. However, the experimental C, H, N and S contents deviated from the theoretical values. This may be due to the drastic hydroscopicity of the charge-transfer complex. It absorbed some water from the air in the dispose process for elemental analysis. Figure 2 may prove this deduction. The pattern of the product adequately exposed to the air only showed the characteristic bands of pyrosulfite, but the pattern of the product from solvent without water showed bands except for the characteristic bands of pyrosulfite. This result indicated that it absorbed some water from the air in the XRD analysis process even without water in the solvent. The distribution of DMI in solid and liquid phases is an important parameter for the absorption of SO2. For this reason, the content of DMI in the liquid phase was determined by gas chromatography and the content of DMI in the solid phase was calculated by mass balance. As shown in Figure 5, the content of DMI in solvent three glycol dimethyl ether decreased quickly once SO2 was introduced into the absorption solution. After 5 minutes, the content of DMI in liquid was only about 30% of the original amount. Twenty minutes later, the conversion of DMI reached 98%, which meant that it had converted completely to a solid. This result indicated that DMI was a benign phase-change absorbent for SO2. 3.3 Further conversion A crystal was prepared by the recrystallization of the absorption product powder in methanol solution. Compared with DMI, the spectrum of the crystal exhibited some new bands at 1002, 1031, 1207 and 1242 cm-1, which can be attributed to the characteristic bands of the S2O62- anion (See Figure 4E)37. Elemental analysis result implied that the product was a dithionate, (C5H9N2)2S2O6, as its experimental C, H, N and S contents of 33.25, 5.01, 15.36, and 18.07% were consistent with the theoretical values (33.89, 5.12, 15.81 and 18.10%). Single crystal XRD analysis further ascertained its structure. As shown in Figure 6(a), a basic structural unit of the crystal consisted of two imidazolium cations and a dithionate anion. The molecule packed through van der Waals force to form a crystal of Pbca space group, and the cell parameters were: a =11.097(2), b = 11.003(2), c = 12.853(3) Å, and α = β = γ = 90° (see Figure 6(b)). Further details

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regarding the structure solution and the final crystal refinements are given in Table 3. Based on the structure of the absorption product, the possible formation chemical equations is proposed as follows: 2(C5H9N2)2S2O5 + O2→ 2(C5H9N2)2S2O6 (3) This result indicated that the absorption product was oxidized by O2 from the air to form dithionate in the recrystallization process. In other words, the absorption product was unstable in air. We have attempted to prepare the single crystal of (C5H9N2)2S2O5 in a vacuum oven. However, the obtained crystal always became oxidized to some degree in the recrystallization process. Thus, the effort to obtain the single crystal of (C5H9N2)2S2O5 was unsuccessful. On the other hand, once the filtrate from the separation process of the solid and liquid was exposed to air for several days, a white solid precipitated from the filtrate. Figure 4F shows its FTIR spectrum. According to the literature38, a free SO4-2 ion under Td symmetry has four fundamental vibrations, the non-degenerate symmetric stretching mode ν1, the doubly degenerate bending mode ν2, the triply asymmetric stretching mode ν3 and the triply degenerate asymmetric bending mode ν4. Only ν3 and ν4 are active in the IR spectrum. The ν3 mode appears as three bands at 1209, 1118 and 1045cm-1 and the ν4 mode is also observed as three bands at 657, 618 and 569 cm-1, as shown in Figure 4F. This result implied that the precipitate was a sulfate. The C, H, N and S contents of the precipitate were found to be 40.78, 6.33, 18.82 and 10.05%, which were consistent with the theoretical values (41.37, 6.25, 19.3, and 11.04%), by taking it as a sulfate (C5H9N2)2SO4 (See Table 2). Single crystal XRD further confirmed its structure. As shown in Figure 7(a), the asymmetric units of precipitate consist of two imidazole cations and an SO42anion. The cation–anion pairs are linked through the N-H…O hydrogen bond. One SO42- anion connected two protonated imidazole cations to form a ribbon along the c-direction, which are interconnected via another ribbon in opposite orientation to form a layer in the bc-plane (See Figure 7(b)). Along the c-direction, SO42- anions are bond to the imidazole cations through single and bifurcated N-H…O hydrogen bonds in a head-to-tail arrangement. The bc layers are linked through the C–H. . .O hydrogen bond to give a three-dimensional network. Stacking of the layers creates a pseudo channel along the a-axis as shown in Figure 7(c). Further details regarding the structure solution and the final crystal refinements are given in Table 3. According to the structure of the absorption product, the possible formation chemical equations is proposed as follows: 4C5H8N2+2H2O+2SO2+O2 → 2(C5H9N2)2SO4 (4) This result indicated that the SO2 in the filtrate would be oxidized slowly to sulfate by oxygen in the air. It should be noted that the stoichiometric rate of the imidazole molecule to the S atom in the sulfate is 2:1, which is higher than that in dithionate (1:1). This result suggested that the formation of sulfate does not take dithionate as an intermediate. 3.4 Thermal stability of the product The thermal stability of (C5H9N2)2S2O5, (C5H9N2)2S2O6 and (C5H9N2)2SO4 was investigated in pure N2 atmosphere. As shown in Figure 8(a), there are three endothermic peaks in the DSC curve of (C5H9N2)2S2O5 at 116, 140 and 164°C. This result indicated that the decomposition of (C5H9N2)2S2O5 occurred through three steps. However, the TG curve indicated that the first and the second weight losses took place almost simultaneously. After that, the residual weight did not exceed 11%, which indicated that the decomposition of pyrosulfite accompanied the volatilization of DMI. In other words, it is difficult to regenerate DMI through traditional heating. The decomposition performance of (C5H9N2)2S2O6 was almost the same to that of (C5H9N2)2S2O5. 6


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However, the residual weight was more than 16% and even the temperature reached 300°C, which was much higher than the boiling point of DMI (204°C). This result indicated that a portion of (C5H9N2)2S2O6 has converted to another kind of compound that must not be DMI. The TG and DSC curves of (C5H9N2)2SO4 are shown in Figure 8(b). The first weight loss began at 50°C, which may be due to the evaporation of water. The highest endothermic peak in the DSC curve appeared at 343°C. The residual weight was about 15% once the temperature was increased to 600°C. This result also indicated that DMI could not be regenerated from the oxidative product (C5H8N2)2SO4 through traditional heating. Traditional method is not suitable for regeneration of DMI. Chemical reaction regeneration may be an alternative way for the recovery of absorption product. The double decomposition reaction of absorption product with base such as Ca(OH)2 or NH3 may be a possible way for the regeneration of DMI. Furthermore, in the work of He et.al39, the SO2 in absorption product was considered as an activated form of SO2, which can be transformed into value-added chemicals by reaction with epoxides. A similar exploration is carrying out in our lab. Conclusions Liquid-solid phase-change absorption of SO2 with imidazoles as the absorbent in organic solvents was investigated in the present work. DMI showed the best phase-change ability among several absorbents. It absorbed SO2 to form pyrosulfite in the presence of water. The absorption product is unstable in air and would be oxidized to dithionate. The absorption rate of DMI for SO2 was very rapid and DMI would almost completely transform to solid phase from liquid phase after 20 minutes. The residual DMI and SO2 in the filtrate would convert to sulfate, (C5H9N2)2SO4, after several days. A TG-DSC experiment indicated that it was difficult to regenerate the absorption or conversion product by traditional heating since they converted to an unknown compound at high temperature. Our work provided some insight for the conversion of SO2 after absorption and gave a possible reason for absorbent inactivation. Acknowledgements The authors acknowledge the financial support from National Natural Science Foundation of China (Grant No. 21666011, 21306071), Yunnan Province Science Foundation (Grant No. 2014FB118), and Analysis and Measure Foundation of Kunming University of Science and Technology (2017M20162108014, 2017M20162208027).

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17. Yang, D.; Hou, M.; Ning, H.; Ma, J.; Kang, X.; Zhang, J.; Han, B., Reversible Capture of SO2 through Functionalized Ionic Liquids. Chemsuschem 2013, 6, (7), 1191-1195. 18. Cui, G. K.; Wang, C. M.; Zheng, J. J.; Guo, Y.; Luo, X. Y.; Li, H. R., Highly efficient SO2 capture by dual functionalized ionic liquids through a combination of chemical and physical absorption. Chemical Communications 2012, 48, (20), 2633-2635. 19. Heldebrant, D. J.; Yonker, C. R.; Jessop, P. G.; Phan, L., Reversible Uptake of COS, CS2, and SO2: Ionic Liquids with O-Alkylxanthate, O-Alkylthiocarbonyl, and O-Alkylsulfite Anions. Chemistry-a European Journal 2009, 15, (31), 7619-7627. 20. Yang, D.; Hou, M.; Ning, H.; Zhang, J.; Ma, J.; Han, B., Efficient SO2 capture by amine functionalized PEG. Physical Chemistry Chemical Physics 2013, 15, (41), 18123-18127. 21. Wang, X. F.; Akhmedov, N. G.; Hopkinson, D.; Hoffman, J.; Duan, Y. H.; Egbebi, A.; Resnik, K.; Li, B. Y., Phase change amino acid salt separates into CO2-rich and CO2-lean phases upon interacting with CO2. Applied Energy 2016, 161, 41-47. 22. Zheng, S. D.; Tao, M. N.; Liu, Q.; Ning, L. Q.; He, Y.; Shi, Y., Capturing CO2 into the Precipitate of a Phase-Changing Solvent after Absorption. Environmental Science & Technology 2014, 48, (15), 8905-8910. 23. Raynal, L.; Bouillon, P. A.; Gomez, A.; Broutin, P., From MEA to demixing solvents and future steps, a roadmap for lowering the cost of post-combustion carbon capture. Chemical Engineering Journal 2011, 171, (3), 742-752. 24. Xu, Z. C.; Wang, S. J.; Chen, C. H., CO2 absorption by biphasic solvents: Mixtures of 1,4-Butanediamine and 2-(Diethylamino)-ethanol. International Journal of Greenhouse Gas Control 2013, 16, 107-115. 25. Pinto, D. D. D.; Zaidy, S. A. H.; Hartono, A.; Svendsen, H. F., Evaluation of a phase change solvent for CO2 capture: Absorption and desorption tests. International Journal of Greenhouse Gas Control 2014, 28, 318-327. 26. Hasib-ur-Rahman, M.; Larachi, F., CO2 Capture in Alkanolamine-RTIL Blends via Carbamate Crystallization: Route to Efficient Regeneration. Environmental Science & Technology 2012, 46, (20), 11443-11450. 27. Zhang, Z.; Zhao, W. B.; Nong, J. J.; Feng, D.; Li, Y. H.; Chen, Y.; Chen, J., Liquid-Solid Phase-Change Behavior of Diethylenetriamine in Nonaqueous Systems for Carbon Dioxide Absorption. Energy Technology 2017, 5, (3), 461-468. 28. Zhao, W.; Zhao, Q.; Zhang, Z.; Liu, J.; Chen, R.; Chen, Y.; Chen, J., Liquid-solid phase-change absorption of acidic gas by polyamine in nonaqueous organic solvent. Fuel 2017, 209, 69-75. 29. Woolven, H.; Gonzalez-Rodriguez, C.; Marco, I.; Thompson, A. L.; Willis, M. C., DABCO-Bis(sulfur dioxide), DABSO, as a Convenient Source of Sulfur Dioxide for Organic Synthesis: Utility in Sulfonamide and Sulfamide Preparation. Organic Letters 2011, 13, (18), 4876-4878. 30. Shannon, M. S.; Irvin, A. C.; Liu, H.; Moon, J. D.; Hindman, M. S.; Turner, C. H.; Bara, J. E., Chemical and Physical Absorption of SO2 by N-Functionalized Imidazoles: Experimental Results and Molecular-level Insight. Industrial & Engineering Chemistry Research 2015, 54, (1), 462-471. 31. Bara, J. E., Versatile and Scalable Method for Producing N-Functionalized Imidazoles. Industrial & Engineering Chemistry Research 2013, 50, (24), 13614-13619. 32. Turner, C. H.; Cooper, A.; Zhang, Z. T.; Shannon, M. S.; Bara, J. E., Molecular Simulation of the Thermophysical Properties of N-Functionalized Alkylimidazoles. Journal of Physical Chemistry B 2012, 116, (22), 6529-6535.

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33. Ando, R. A.; Matazo, D. R. C.; Santos, P. S., Detailed analysis of the charge transfer complex N,N-dimethylaniline-SO2 by Raman spectroscopy and density functional theory calculations. Journal of Raman Spectroscopy 2010, 41, (7), 771-775. 34. Faria, D.; Santos, P., Raman and infrared spectra of some aromatic amine–sulphur dioxide molecular complexes. Journal of Raman spectroscopy 1988, 19, (7), 471-478. 35. Oh, J. J.; Labarge, M. S.; Matos, J.; Kampf, J. W.; Ii, K. W. H.; Kuczkowski, R. L., Structure of the trimethylamine-sulfur dioxide complex. Journal of Physical Chemistry 1991, 95, (19), 4732-4738. 36. Devarajan, V.; Shurvell, H. F., Vibrational spectra and normal coordinate analysis of crystalline potassium pyrosulfite K 2 S 2 O 5. Spectrochimica Acta Part A Molecular Spectroscopy 1977, 33, (11), 1041-1047. 37. Beattie, I. R.; Gall, M. J.; Ozin, G. A., Single-crystal Raman studies and the vibrational spectrum of the dithionate ion. Journal of the Chemical Society A Inorganic Physical Theoretical 1969, 1001-1008. 38. Guerfel, T.; Jouini, A., Crystal structure, thermal analysis, and IR spectroscopic investigation of bis(2-amino-6-methyl) pyridinium sulfate. Journal of Chemical Crystallography 2005, 35, (7), 513-521. 39. Yang, Z. Z.; He, L. N.; Zhao, Y. N.; Yu, B., Highly Efficient SO2 Absorption and Its Subsequent Utilization by Weak Base/Polyethylene Glycol Binary System. Environmental Science & Technology 2013, 47, (3), 1598-1605.

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Table 1. Liquid-solid phase-change absorption of SO2 by absorbent in organic solvent (wt.30%) solvent

propylene

three glycol butanol

absorbent

dimethylacetamide

carbonate

1,3,5-trimethylpyraz

dimethyl ether

No change

No change

No change

No change

2-methylimidazole

undissolved

No change

No change

undissolved

1-methylimidazole

No change

No change

No change

white solid

DMI

white solid

white solid

white solid

ole

a little white solid

Table 2. Mass fraction C, H, N and S in precipitate Reaction

Chemical formula

DMI,SO2 and H2O

[(C5H9N2)2S2O5]

DMI, SO2, O2 and H2O in methanol

[(C5H9N2)2S2O6]

DMI, SO2, O2 and H2O in filtrate

[(C5H9N2)2SO4]

Results

C

H

N

S

Theoretical value

35.49

5.36

16.56

18.85

Experimental value

35.76

5.32

16.84

18.17

Theoretical value

33.89

5.12

15.81

18.10

Experimental value

33.25

5.01

15.36

18.07

Theoretical value

41.37

6.25

19.30

11.04

Experimental value

40.78

6.33

18.82

10.05

Table 3. Crystal data and structure refinement for (C5H9N2)2S2O6 and (C5H9N2)2SO4 Parameters

(C5H9N2)2S2O6

(C5H9N2)2SO4

Empirical formula

C10H18N4O6S2

C10H18N4O4S

Formula weight (g/mol)

356.42

290.34

Crystal system

orthorhombic

Monoclinic

space group

Pbca

P2(1)

a (Å)

11.097(2)

7.1542(14)

b (Å)

11.003(2)

13.789(3)

c (Å)

12.853(3)

7.3260(15)

alpha (deg)

90.000

90.000

beta (deg)

90.000

109.95(3)

gamma (deg)

90.000

90.000

Volume (Å^3)

1567.16(50)

679.3(2)

Z

4

2

Calculated density (g/cm^3)

1.1054

1.419

Absorption coefficient (mm^-1)

0.37

0.255

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Energy & Fuels

Theta range for data collection (deg)

1.51 to 51.99 deg.

3.03 to 26.00 deg.

Limiting indices

-13=< h =< 12, -13=< k =< 13,

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Graphic Abstract - American Chemical Society

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