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Reversible absorption of SO2 by amino acid aqueous solutions
Journal of Hazardous Materials 229–230 (2012) 398–403
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Journal of Hazardous Materials journal homepage: www.elsevier.com/locate/jhazmat
Reversible absorption of SO2 by amino acid aqueous solutions Renpan Deng, Lishan Jia ∗ , Qianqian Song, Shuai Su, Zhongbiao Tian Department of Chemical Engineering and Biochemical Engineering, College of Chemistry and Chemical Engineering, Xiamen University, Xiamen 361005, China
h i g h l i g h t s
-Alanine solution is found to be excellent absorbent for SO2 removal. 20–30 ◦ C is optimal for SO2 absorption, and 150 ◦ C is optimal for desorption. The neutral environment (pH = 6.8) was found to be optimal for SO2 removal. SO2 interacts with -alanine by some weak interactions, such as hydrogen bonds. The absorbent has an excellent regeneration performance.
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Article history: Received 21 February 2012 Received in revised form 10 June 2012 Accepted 11 June 2012 Available online 20 June 2012 Keywords: Amino acid Absorption -Alanine Sulfur dioxide Flue gas desulfurization
a b s t r a c t Six water-soluble amino acids (glycine, l-␣-alanine, dl-alanine, -alanine, proline and arginine) aqueous solutions were applied to remove SO2 from SO2 –N2 system in this report. All the tested amino acids solutions were found to be excellent absorbents for SO2 removal, and SO2 saturation uptake of -alanine solution was the highest under the same experimental conditions. The effects of amino acid concentration, SO2 concentration, absorption temperature, desorption temperature and initial pH value of the absorbent on the removal of SO2 were investigated with -Ala solution. The experimental results showed that SO2 saturation uptake increased with the increase in -alanine solution and SO2 concentration. Room temperature (20–30 ◦ C) was found to be optimal for SO2 absorption. Additionally the SO2 desorption capacity increased with increasing desorption temperature. The neutral environment pH value of 6.8 was found to be optimal for SO2 removal. Ten continuous absorption–desorption cycles showed that the absorbent had an excellent regeneration performance. 13 C NMR and ultraviolet analyses offer ample evidence to speculate that the bonding between SO2 and -alanine was not covalent but some weak interactive forces, such as dispersion force, induction force, dipole–dipole force and hydrogen bond. © 2012 Elsevier B.V. All rights reserved.
1. Introduction Sulfur in coal and oil is converted to sulfur dioxide (SO2 ) during combustion. SO2 is a conspicuous atmospheric pollutant that poses major threats to the environment and human life [1–4]. Currently, flue gas desulfurization (FGD) is one of the most widely used techniques to control SO2 emission [5]. However, the main flaw of the technique is the difficult regeneration of absorbents [6,7]. In addition, although SO2 is a very important and useful source for many intermediates in chemical productions [8], it is difficult to recycle SO2 during the FGD process as it forms stable sulfite or sulfate. Consequently, it is particularly important to develop renewable absorbents for SO2 removal. Several absorbents have been developed over the years; these can be categorized as inorganic salts, organic solvents and ionic
∗ Corresponding author. Tel.: +86 592 2188283; fax: +86 592 2184822. E-mail address: [email protected] (L. Jia). 0304-3894/$ – see front matter © 2012 Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.jhazmat.2012.06.020
liquids (ILs). Among the inorganic salts, the use of Na2 CO3 and NaHCO3 solutions for desulfurization has been widely explored [9–11]. The desulfurization principle of such absorbents is similar to that of limestone and lime in that it leads to the formation of stable sulfites or sulfates resulting in difficult regeneration of the absorbents. Organic solvents have been used for SO2 removal since the forties of last century [12]. Many organic solvents, such as dimethylacetamide [13], methyldiethanolamine [14] and ethylenediamine [15] have been found to be good solvents for SO2 . However, these solvents are highly toxic, volatile, degradable and as such not environmentally friendly [16,17]. Ionic liquids (ILs) have attracted much attention from the industrial and academic communities as ‘green solvents’ due to their remarkable properties: extremely low vapor pressure, high thermal and chemical stability, and excellent solvating power for both organic and inorganic compounds [18–20]. Wu and co-workers earlier on reported that 1,1,3,3-tetramethylguanidine (TMG) lactate ILs could effectively absorb SO2 from simulated flue gas [21]. The use of other ILs such as those based on 1,1,3,3-tetramethylguanidinium
R. Deng et al. / Journal of Hazardous Materials 229–230 (2012) 398–403
399
Table 1 Some basic information of the six tested amino acids. Appellation
Abbreviation
Molecular weight (g/mol)
Molecular structure
O Glycine
Gly
H2N
75
CH C
OH
H O l-␣-Alanine
l-␣-Ala
H2N
89
CH C
OH
CH3
H dl-Alanine
dl-Ala
H3C
89
H COOH HOOC
C
C
CH3
NH2
NH2
O -Alanine
-Ala
Proline
Pro
89
H2N
H2C
H N
C
115
CH2 C
OH
O
OH
H2N Arginine
Arg
174
C NH
H N
NH2 H2 H2 H2 C C C CH C
O
OH
[22] and imidazolium [23] for the removal of SO2 have also been explored. However, commercial ILs are expensive and their preparation processes are complicated resulting in much higher costs and limiting their industrial applications. Recently, we demonstrated that l-␣-alanine supported on Al2 O3 could remove SO2 efficiently and effectively [24]. In this present paper, we have further expanded our work, and found that several water-soluble amino acid (Gly, l-␣-Ala, dl-Ala, -Ala, Pro and Arg) aqueous solutions showed excellent absorption performances on SO2 , such as high absorption capacity and good regeneration performance. So these amino acids could be alternative desulfurization absorbents. In fact, amino acids as biomolecules are the building blocks for peptides and proteins [25], and are usually used in biochemical and biological syntheses. To our knowledge, amino acids are seldom directly used for environment-related issues. As an extension herein, we have studied the desulfurization performance of six water-soluble amino acids (Gly, l-␣-Ala, dl-Ala, -Ala, Pro and Arg) aqueous solutions. Some basic information of these amino acids is presented in Table 1. The effects of the various factors on the desulfurization performance such as amino acid concentration, SO2 concentration, absorption and desorption temperature, and the initial pH value of absorbent are investigated by using -Ala solution. It is hoped that this work could enlighten researchers on the development of such desulfurization absorbents.
relatively long, for two specific advantages: (1) it prevents entrainment caused by the relatively high gas flow rate and (2) it facilitates refluxing and condensing during the desorption process (it requires as high as 150 ◦ C). Approximately 5.0 mL amino acid solution of appropriate concentration was placed in the tube reactor to remove SO2 from simulated flue gas. The simulated flue gas was a mixture of SO2 and N2 with appropriate concentration. The flow rate of the mixture gas was adjusted by a mass flow meter to 60 mL/min. The SO2 concentration changes at the outlet of the reactor were determined by iodine titration method (HJ/T 56-2000, a standard method of State Environmental Protection Administration of China). The amount of SO2 absorption capacity was expressed as milligram SO2 per milliliter amino acid solution (mg/mL). 2.2. SO2 desorption experiments After the saturation absorption (SO2 concentration at the outlet of the reactor was kept constant), the gas mixture was switched to N2 and the temperature increased to the required desorption temperature at a rate of 10 ◦ C/min; and the N2 flow was maintained until SO2 was no longer detected in the vent. Again, the concentration of SO2 at outlet of the reactor was determined by iodine titration method and the amount of desorbed SO2 expressed as before (mg/mL). Meanwhile, the absorbent was regenerated during the SO2 desorption process.
2. Materials and methods 2.3. 2.1. SO2 absorption experiments All the amino acids (biochemical reagent (BR)) used in the experiments were purchased from Sinopharm Chemical Reagent Co., Ltd, China. Experiments were performed in a vertically placed glass-made tube reactor (with 8 mm inside diameter and 39.5 cm long) under isothermal conditions. The tube reactor was made
13 C
NMR spectra
The original -Ala solution, -Ala solution saturated with SO2 , -Ala solution after SO2 desorption at 150 ◦ C and -Ala solution after ten continuous absorption–desorption cycles were mixed with heavy water (D2 O) in a volume ratio of 9:1 respectively. 13 C spectra were then recorded on a Varian Unity INOVA 100 MHz instrument.
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2.4. Ultraviolet spectra Ultraviolet (UV) spectra was use to monitor the original -Ala solution, the -Ala solution saturated with SO2 , -Ala solution after SO2 desorption at 150 ◦ C and the water saturated with SO2 . Samples of the solutions were diluted to the appropriate concentrations with deionized water and subsequently used for detection analyses. UV spectroscopy detection of the samples was carried out on TU-1900 (Purkinje General Instrument Co., Ltd., Beijing) spectrophotometers at a resolution of 1 nm. The spectra were recorded in the region of 190–400 nm. 3. Results and discussion 3.1. Activity of different amino acid solutions on SO2 removal Fig. 1a illustrates the activity of different amino acid (Gly, l-␣Ala, dl-Ala, -Ala, Pro and Arg) aqueous solutions on the removal of SO2 at 30 ◦ C. All the tested amino acid solutions exhibited good activity on the removal of SO2 compared with water. SO2 saturation uptake by -Ala solution was higher than that of the other ␣-amino acid solutions. This suggests that the structure of amino acids influences SO2 absorption to a great extent. In comparison to some other absorbents mentioned previously, for example, TMG ILs (SO2 saturation uptake is 305 mg/g, absorbed at 40 ◦ C, gas composition (volume fraction (vol.%): 8% SO2 and 92% N2 [21])) and imidazolium-based ILs (SO2 saturation uptake is 20.2 mg/g, absorbed at 20 ◦ C, gas composition: 10% SO2 and 90% N2 [23]), it is worth mentioning that SO2 saturation uptake by 1.0 M -Ala solution is relatively higher (about 390 mg/g). Fig. 1b shows that SO2 saturation uptake increases linearly with the increase in -Ala concentration. Increasing the concentration by 0.5 M increases the saturation uptake by an average of about 19.5 mg/mL. 3.2. Effects of SO2 concentration on SO2 removal by ˇ-Ala solution In Fig. 2, the effects of SO2 concentration on SO2 removal by -Ala solution at 30 ◦ C is shown. Fig. 2a displays the relationship between the rate of absorption and SO2 concentration. C0 represents the initial concentration of SO2 and Ct represents the concentration of SO2 at the exit of the reactor at time t. Ct /C0 could reflect the absorption rate in here. The saturation time (the time the absorption reaches saturation, i.e. at Ct /C0 = 1) decreased with the increase in the SO2 concentration. This could be attributed to the diffusion driving force increased with the increase in SO2 concentration, which increased SO2 absorption rate and thus
Fig. 1. (a) The activity of different amino acid aqueous solutions on SO2 removal. Absorption temperature: 30 ◦ C; concentration of amino acids: 1.0 M; gas composition (vol.%), SO2 : 4.1%; balance gas: N2 . (b) The activity of -Ala solution with different concentration on SO2 removal. Absorption temperature: 30 ◦ C; gas composition (vol.%): SO2 , 4.1%; balance gas: N2 .
promoted SO2 removal. Therefore, SO2 saturation uptake increased with the increasing of SO2 concentration as shown in Fig. 2b. However, a saturation uptake of about 26.8 mg/mL was realized when the concentration of SO2 was decreased to 0.8% revealing that the saturation uptake of SO2 decreased only by about 33% when the concentration of SO2 was decreased by about five times. This demonstrates that the -Ala solution could be employed to remove SO2 under relatively low SO2 concentrations.
Fig. 2. Effects of SO2 concentration on SO2 removal by -Ala solution. (a) Breakthrough curve, (b) variation of SO2 saturation uptake with SO2 concentration, absorption temperature: 30 ◦ C; concentration of -Ala (C-Ala ) is 1.0 M; balance gas: N2 .
R. Deng et al. / Journal of Hazardous Materials 229–230 (2012) 398–403
Fig. 3. Effects of temperature on SO2 removal by -Ala solution. C-Ala = 1.0 M; gas composition (vol.%): SO2 , 4.1%; balance gas: N2 .
3.3. Effects of absorption temperature on SO2 removal by ˇ-Ala solution Temperature is one of the main factors that affect the absorption of SO2 . Fig. 3 shows that SO2 saturation uptake was basically constant between 20 and 30 ◦ C but decreased (about 20.7%) with further increase in the absorption temperature to 60 ◦ C. This result clearly indicates that high absorption temperature was not favorable for SO2 removal, and as such room temperature range (20–30 ◦ C) was found to be the optimal condition for SO2 removal. This could be explained by the fact that high temperatures are not conducive for gaseous dissolution.
401
Fig. 5. Effects of initial pH value on SO2 removal by -Ala solution. Absorption temperature: 30 ◦ C; C-Ala = 1.0 M; gas composition (vol.%): SO2 , 4.1%; balance gas: N2 .
simply be due to SO2 physically dissolving in -Ala solution, so that it could be desorbed at the relatively low temperature of 30 ◦ C. Further still, SO2 desorption capacity and desorption rate are observed to respectively increased with increase in desorption temperature, and at a temperature of 150 ◦ C, SO2 had been desorbed almost completely. This result suggests that SO2 interacts with -Ala by different mechanisms besides by the simple physical dissolution. The SO2 dissolution by -Ala solution by different mechanisms requires different energy quantities for desorption and therefore, SO2 capacity increased with the increase in desorption temperature.
3.4. Effects of desorption temperature on SO2 desorption capacity
3.5. Effects of initial pH value on SO2 removal by ˇ-Ala solution
Fig. 4 reveals the relation between desorption temperature and SO2 desorption capacity. The SO2 desorption process was carried out after the SO2 absorption reached its saturation at 30 ◦ C. The processes were then operated at temperatures of 30, 60, 90, 120 and 150 ◦ C for 150 min under N2 flow. SO2 desorption capacity and desorption rate ((desorption capacity/saturation uptake) × 100%) were about 10.1 mg/mL and 25.2%, respectively, at the desorption temperature of 30 ◦ C. This implied that only a small portion of SO2 was desorbed at this temperature. The easily desorbed SO2 could
The initial pH value of -Ala solution was adjusted to 0.9, 2.9, 5.0 and 10.1, respectively, by addition of hydrochloric acid solution (1 M) or sodium hydroxide solution (1 M). The pH value of neutral -Ala solution (without any added acids or bases) was 6.8. Fig. 5 shows that SO2 saturation uptake is low when the adjusted pH of -Ala solution is low (4.5 mg/mL at 0.9 and 5.5 mg/mL at 2.9) due possibly, to the large H+ in the solution inhibiting the dissolution of SO2 . On the contrary, SO2 saturation uptake is higher with high adjusted pH (61.1 mg/mL at 10.1), even higher than that at the neutral pH (40.1 mg/mL at 6.8). However, desorption of SO2 was incomplete with desorption rate of 73.4% due partly to reaction SO2 with NaOH to form stable Na2 SO3 or NaHSO3 . Thus acidic -Ala solution does not promote SO2 absorption and alkaline Ala solution is not conducive for SO2 desorption. The neutral -Ala solution (pH = 6.8) is the optimal pH for SO2 removal. Amino acids (RCHNH2 COOH) in aqueous solution usually exist in the form of zwitterion (RCHNH3 + COO− ), and it may react with either H+ or OH− as illustrated in the example with -Ala reaction equation below: NH+ CHCH2 COO− + H+ NH+ CHCH2 COOH 3 3 CHCH2 COO− + OH− NH2 CHCH2 COO− + H2 O NH+ 3 Therefore, the absorption of SO2 in the neutral -Ala solution (pH = 6.8) may be described as below: 1. Dissolution of SO2 in water,
Fig. 4. Effects of desorption temperature on SO2 desorption capacity. Absorption temperature: 30 ◦ C, C-Ala = 1.0 M; gas composition (vol.%): SO2 , 4.1%; balance gas: N2 .
SO2 + H2 O H+ + HSO− 3
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Fig. 6. Regeneration performance of -Ala solution on SO2 removal. Absorption temperature: 30 ◦ C, C-Ala = 1.0 M; gas composition (vol.%): SO2 , 4.1%; balance gas: N2 .
2. Combination of carboxyl (from -Ala) with H+ , NH+ CH2 CH2 COO− 3
+
+H
NH+ CH2 CH2 COOH 3
This reaction greatly promotes the dissolution of SO2 . + 3. HSO− 3 further reacted with NH3 CH2 CH2 COOH, and this step will be discussed in detail in Section 3.7.
Fig. 7. 13 C NMR spectra (100 MHz) of original -Ala solution (1), -Ala solution saturation with SO2 (2), -Ala solution after SO2 desorption at 150 ◦ C (3) and Ala solution after ten continuous absorption–desorption cycles at the desorption temperature of 150 ◦ C (4). (Absorption temperature: 30 ◦ C; C-Ala = 1.0 M; gas composition (vol.%): SO2 , 4.1%; balance gas: N2 .)
solution was reversible and that -Ala was regenerated after SO2 desorption. This is consistent with the obtained experimental data.
3.6. Regeneration performance of ˇ-Ala solution The regeneration experiments were carried out at 30, 120 and 150 ◦ C desorption temperatures. Results of ten continuous absorption-desorption cycles of the -Ala solution are shown in Fig. 6, depicting constant SO2 desorption capacity at each temperature as a result of excellent regenerative ability. The indication is that the interaction between the SO2 and the -Ala solution is a reversible equilibrium reaction. 3.7. Absorption mechanism 3.7.1. 13 C NMR analysis The 13 C NMR analysis was used to explore the SO2 absorption mechanism. The 13 C NMR spectra of original -Ala solution (spectra 1), -Ala solution saturation with SO2 (spectra 2), Ala solution after SO2 desorption at 150 ◦ C (spectra 3) and -Ala solution after ten continuous absorption–desorption cycles at the desorption temperature of 150 ◦ C (spectra 4) are shown in Fig. 7. Three resonances observed at ı = 19.8, 22.9 and 164.5 ppm in original -Ala solution (NH3 + CH2 CH2 COO− ) are assigned to the carbon of CH2 , CH2 NH3 + and COO− respectively which shift to 17.5, 21.6 and 161.1 ppm respectively after saturation with SO2 . This chemical shift could be attributed to the result of SO2 interacting with -Ala. The results imply that there was no covalent bond formation between SO2 and -Ala, otherwise much bigger chemical shift would have been observed. The resonance of COO− changed by 3.4 ppm (from 164.5 to 161.1 ppm) after -Ala solution saturation with SO2 , while the resonances of CH2 and NH3 + CH2 changed by 2.3 and 1.3 ppm, respectively. This suggests that the SO2 interacted with COO− , and as a result, the chemical shift of COO− , CH2 and CH2 NH3 + showed a decreasing trend. All the three resonances were essentially restored to their original states after SO2 desorption at 150 ◦ C (spectra 3). Furthermore, the spectra of -Ala solution got restored after ten continuous absorption–desorption cycles (spectra 4). This result suggests that the absorption of SO2 by -Ala
3.7.2. UV spectra analysis Fig. 8 depicts the ultraviolet spectra of original -Ala solution (spectra 1), -Ala solution saturation with SO2 (spectra 2), -Ala solution after SO2 desorption at 150 ◦ C (spectra 3) and water saturation with SO2 (spectra 4). As mentioned earlier (Section 3.5), SO2 dissolves in water and ionized partly to form bisulfite ion. The two peaks observed at 276 nm and 204 nm in spectra 4 are assigned to SO2 and HSO3 − [26,27], respectively. The absorption peak of SO2 at 276 nm is also observed in the ultraviolet spectra of the -Ala solution saturated with SO2 (spectra 2) confirming that small portion of the SO2 indeed simply dissolves physically in the -Ala solution, as noted in Section 3.4. Comparatively, one notable difference between the ultraviolet spectra of water saturated with SO2 (spectra 4), and the spectra of -Ala solution saturated with
Fig. 8. Ultraviolet spectra of original -Ala solution (1), -Ala solution saturation with SO2 (2), -Ala solution after SO2 desorption at 150 ◦ C (3) and water saturation with SO2 (4). (Absorption temperature: 30 ◦ C, C-Ala = 1.0 M; gas composition (vol.%): SO2 : 4.1%; balance gas: N2 .).
R. Deng et al. / Journal of Hazardous Materials 229–230 (2012) 398–403
SO2 (spectra 2) was that the absorption peak of HSO3 − at 204 nm shifted to 210 nm. The absorption peak at 214 nm in the spectra of the original -Ala solution (spectra 1) is assigned to -Ala. This peak also shifted slightly after the -Ala solution is saturated with SO2 (spectra 2) because of the interaction of HSO3 − with -Ala. After the SO2 desorption at 150 ◦ C (spectra 3), the absorption peak (of SO2 ) at 276 nm disappeared and that of -Ala is restored (to 214 nm). Note that the -Ala solution saturated with SO2 was diluted 50 times before the UV scanning (otherwise its absorbance would have been beyond the instrument’s measurable range), while the rest of the other solutions were diluted 10 times. This suggests that the interaction between SO2 and -Ala has a hyperchromic effect on the UV absorbance of -Ala. Consequently, we speculate that SO2 dissolves in -Ala solution with a high saturation uptake because the interaction between them is not only through the simple physical dissolution; and that the interaction between them is not by covalent bonding but rather some weak forces of interactions, such as dispersion force, induction force, dipole–dipole force and hydrogen bonding. This observation is in agreement with that of van Dam who used organic solvents to remove SO2 [14]. In this study, we have shown that SO2 interacts with the carboxyl of -Ala to form hydrogen bond through a two-step process; the first step involves the dissolution of SO2 in water that is then ionized to generate hydrogen and bisulfite ions and the hydrogen ions combine with COO− to form COOH as mentioned earlier (Section 3.5). In the second step, hydrogen sulfite ions interact with COOH through hydrogen bonding. The process is summarized as:
In the above reaction with the dotted line representing hydrogen bond, the SO2 can be desorbed by heating. This partly explains why the SO2 saturation uptake by -Ala was higher than that of the other tested ␣-amino acids (Section 3.1). NH3 + is electron-withdrawing group, and has an electron withdrawing inductive effect on COOH. ␣-NH3 + is close to COOH and as such the inductive effect of electron withdrawing is much stronger in this case. This leads to the electrons being deflected from the COOH to NH3 + , which is not conducive for the formation of the hydrogen bond. On the contrary, -NH3 + is far from COOH, so it has little influence on the COOH group and therefore has little effect on the formation of hydrogen bond. In a nutshell, SO2 easily interacts with -Ala to form hydrogen bonds and therefore, the SO2 saturation uptake of -Ala is much higher than that observed for the other tested ␣-amino acids.
4. Conclusion Six water-soluble amino acids (Gly, l-␣-Ala, dl-Ala, -Ala, Pro and Arg) solutions were employed to remove SO2 from SO2 –N2 system. The effect of concentrations of the -Ala solution and SO2 , absorption and desorption temperatures and the initial pH of the -Ala solution on the desulfurization performance were investigated. Based on the results we conclude that: (1) The amino acid solutions have good activity on SO2 removal with -Ala solution exhibiting the highest saturation uptake. In fact, its saturation uptake increases with increasing concentrations of SO2 and -Ala solution. (2) While the optimal removal of SO2 is at room temperature (20–30 ◦ C), its optimal desorption is at high temperature (150 ◦ C). (3) Though acidic solution (pH = 0.9; 2.9) is not conducive for SO2 absorption, alkaline solution (pH = 10.1) is not conducive for SO2 desorption either; however, neutral solution (pH = 6.8) offers the optimal pH value for SO2 removal. (4) The interaction between SO2 and -Ala is not by covalent bonding but rather some weak interactive forces (such as dispersive, inductive, dipole-dipole
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and hydrogen bonding). (5) The absorbent (-Ala solution) posses excellent regenerative ability. Acknowledgements This project was supported by the Fujian Province Science & Technology Program of China (2010H6024). The authors are grateful for the State Key Laboratory for Physical Chemistry of Solid Surface of Xiamen University and Mr. Ibrahim Abdul-Rauf and Fazak Siyekwo for revision of the paper. References [1] M.J. Molina, L.T. Molina, Megacities and atmospheric pollution, J. Air Waste Manage. Assoc. 54 (2004) 644–680. [2] J.F. Izquierdo, C. Fite, F. Cunill, M. Iborra, J. Tejero, Kinetic study of the reaction between sulfur dioxide and calcium hydroxide at low temperature in a fixedbed reactor, J. Hazard. Mater. B76 (2000) 113–123. [3] K.T. Lee, A.R. Mohamed, S. Bhatia, K.H. Chu, Removal of sulfur dioxide by fly ash/CaO/CaSO4 sorbents, Chem. Eng. J. 114 (2005) 171–177. [4] I. Dahlan, K.T. Lee, A.H. Kamaruddin, A.R. Mohamed, Evaluation of various additives on the preparation of rice husk ash (RHA)/CaO-based sorbent for flue gas desulfurization (FGD) at low temperature, J. Hazard. Mater. 161 (2009) 570–574. [5] F.J. Gutiérrez Ortiz, A simple realistic modeling of full-scale wet limestone FGD units, Chem. Eng. J. 165 (2010) 426–439. [6] G.L. Hu, D.J. Kim, S. Wedel, J.P. Hansen, Review of the direct sulfation reaction of limestone, Prog. Energy Combust. 32 (2006) 386–407. [7] R.K. Srivastava, W. Jozewicz, C. Singer, SO2 scrubbing technologies: a review, Environ. Prog. 20 (2001) 219–228. [8] J. Huang, A. Riisager, P. Wasserscheid, R. Fehrmanna, Reversible physical absorption of SO2 by ionic liquids, Chem. Commun. 38 (2006) 4027–4029. [9] K. Stejskalova, I. Spirovová, E. Lippert, K. Mocek, Z. Bastl, A study of SO2 interaction with sodium carbonates by X-ray photoelectron spectroscopy, Appl. Surf. Sci. 103 (1996) 509–516. [10] C.F. Wu, S.J. Khang, T.C. Keener, S.K. Lee, A model for dry sodium bicarbonate duct injection flue gas desulfurization, Adv. Environ. Res. 8 (2004) 655–666. [11] S. Ebrahimi, C. Picioreanu, R. Kleerebezem, J.J. Heijnen, M.C.M. van Loosdrecht, Rate-based modelling of SO2 absorption into aqueous NaHCO3 /Na2 CO3 solutions accompanied by the desorption of CO2 , Chem. Eng. Sci. 58 (2003) 3589–3600. [12] G.H. Gleason, N.J. Montclair, A.C. Loonam, Recovery of sulphur dioxide, U.S. patent 21.106.453, 1940. [13] R.J. Demyanovich, S. Lynn, Prediction of infinite dilution activity coefficients of sulfur dioxide in organic solvents, J. Solution Chem. 20 (1991) 693–701. [14] M.H.H. van Dam, A.S. Lamine, D. Roizard, P. Lochon, C. Roizard, Selective sulfur dioxide removal using organic solvents, Ind. Eng. Chem. Res. 36 (1997) 4628–4637. [15] Z.G. Tang, C.C. Zhou, C. Chen, Studies on flue gas desulfurization by chemical absorption using an ethylenediamine–phosphoric acid solution, Ind. Eng. Chem. Res. 43 (2004) 6714–6722. [16] I.J. Uyanga, O.I. Raphael, Studies of SO2 and O2 induced degradation of aqueous MEA during CO2 capture from power plant flue gas streams, Ind. Eng. Chem. Res. 46 (2007) 2558–2566. [17] T. Supap, I. Raphael, P. Tontiwachwuthikul, C. Saiwan, Kinetics of sulfur dioxideand oxygen-induced degradation of aqueous monoethanolamine solution during CO2 absorption from power plant flue gas streams, Int. J. Green. Gas Con. 3 (2009) 133–142. [18] K. Fukumoto, M. Yoshizawa, H. Ohno, Room temperature ionic liquids from 20 natural amino acids, J. Am. Chem. Soc. 127 (2005) 2398–2399. [19] T. Welton, Room-temperature ionic liquids. Solvents for synthesis and catalysis, Chem. Rev. 99 (1999) 2071–2083. [20] Y.Y. Jiang, G.N. Wang, Z. Zhou, Y.T. Wu, J. Geng, Z.B. Zhang, Tetraalkylammonium amino acids as functionalized ionic liquids of low viscosity, Chem. Commun. 4 (2008) 505–507. [21] W.Z. Wu, B.X. Han, H.X. Gao, Z.M. Liu, T. Jiang, J. Huang, Desulfurization of flue gas: SO2 absorption by an ionic liquid, Angew. Chem. Int. Ed. 43 (2004) 2415–2417. [22] J. Huang, A. Riisager, R.W. Berg, R. Fehrmann, Tuning ionic liquids for high gas solubility and reversible gas sorption, J. Mol. Catal. A: Chem. 279 (2008) 170–176. [23] J. Huang, A. Riisager, P. Wasserscheid, R. Fehrmann, Reversible physical absorption of SO2 by ionic liquids, Chem. Commun. 38 (2006) 4027–4029. [24] R.P. Deng, L.S. Jia, Reversible removal of SO2 at low temperature by l-␣-alanine supported on ␥-Al2 O3 , Fuel 93 (2012) 385–390. [25] Z. Paˇı szti, L. Guczi, Amino acid adsorption on hydrophilic TiO2 : a sum frequency generation vibrational spectroscopy study, Vib. Spectrosc. 50 (2009) 48–56. [26] K. Fuwa, Bert L. Vallee, Molecular flame absorption spectrometry for sulfur, Anal. Chem. 41 (1969) 188–190. [27] A. Syty, Determination of sulfur dioxide by ultraviolet absorption spectrometry, Anal. Chem. 45 (1973) 1744–1747.
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Reversible absorption of SO2 by amino acid aqueous solutions Renpan Deng, Lishan Jia ∗ , Qianqian Song, Shuai Su, Zhongbiao Tian Department of Chemical Engineering and Biochemical Engineering, College of Chemistry and Chemical Engineering, Xiamen University, Xiamen 361005, China
h i g h l i g h t s
-Alanine solution is found to be excellent absorbent for SO2 removal. 20–30 ◦ C is optimal for SO2 absorption, and 150 ◦ C is optimal for desorption. The neutral environment (pH = 6.8) was found to be optimal for SO2 removal. SO2 interacts with -alanine by some weak interactions, such as hydrogen bonds. The absorbent has an excellent regeneration performance.
a r t i c l e
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Article history: Received 21 February 2012 Received in revised form 10 June 2012 Accepted 11 June 2012 Available online 20 June 2012 Keywords: Amino acid Absorption -Alanine Sulfur dioxide Flue gas desulfurization
a b s t r a c t Six water-soluble amino acids (glycine, l-␣-alanine, dl-alanine, -alanine, proline and arginine) aqueous solutions were applied to remove SO2 from SO2 –N2 system in this report. All the tested amino acids solutions were found to be excellent absorbents for SO2 removal, and SO2 saturation uptake of -alanine solution was the highest under the same experimental conditions. The effects of amino acid concentration, SO2 concentration, absorption temperature, desorption temperature and initial pH value of the absorbent on the removal of SO2 were investigated with -Ala solution. The experimental results showed that SO2 saturation uptake increased with the increase in -alanine solution and SO2 concentration. Room temperature (20–30 ◦ C) was found to be optimal for SO2 absorption. Additionally the SO2 desorption capacity increased with increasing desorption temperature. The neutral environment pH value of 6.8 was found to be optimal for SO2 removal. Ten continuous absorption–desorption cycles showed that the absorbent had an excellent regeneration performance. 13 C NMR and ultraviolet analyses offer ample evidence to speculate that the bonding between SO2 and -alanine was not covalent but some weak interactive forces, such as dispersion force, induction force, dipole–dipole force and hydrogen bond. © 2012 Elsevier B.V. All rights reserved.
1. Introduction Sulfur in coal and oil is converted to sulfur dioxide (SO2 ) during combustion. SO2 is a conspicuous atmospheric pollutant that poses major threats to the environment and human life [1–4]. Currently, flue gas desulfurization (FGD) is one of the most widely used techniques to control SO2 emission [5]. However, the main flaw of the technique is the difficult regeneration of absorbents [6,7]. In addition, although SO2 is a very important and useful source for many intermediates in chemical productions [8], it is difficult to recycle SO2 during the FGD process as it forms stable sulfite or sulfate. Consequently, it is particularly important to develop renewable absorbents for SO2 removal. Several absorbents have been developed over the years; these can be categorized as inorganic salts, organic solvents and ionic
∗ Corresponding author. Tel.: +86 592 2188283; fax: +86 592 2184822. E-mail address: [email protected] (L. Jia). 0304-3894/$ – see front matter © 2012 Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.jhazmat.2012.06.020
liquids (ILs). Among the inorganic salts, the use of Na2 CO3 and NaHCO3 solutions for desulfurization has been widely explored [9–11]. The desulfurization principle of such absorbents is similar to that of limestone and lime in that it leads to the formation of stable sulfites or sulfates resulting in difficult regeneration of the absorbents. Organic solvents have been used for SO2 removal since the forties of last century [12]. Many organic solvents, such as dimethylacetamide [13], methyldiethanolamine [14] and ethylenediamine [15] have been found to be good solvents for SO2 . However, these solvents are highly toxic, volatile, degradable and as such not environmentally friendly [16,17]. Ionic liquids (ILs) have attracted much attention from the industrial and academic communities as ‘green solvents’ due to their remarkable properties: extremely low vapor pressure, high thermal and chemical stability, and excellent solvating power for both organic and inorganic compounds [18–20]. Wu and co-workers earlier on reported that 1,1,3,3-tetramethylguanidine (TMG) lactate ILs could effectively absorb SO2 from simulated flue gas [21]. The use of other ILs such as those based on 1,1,3,3-tetramethylguanidinium
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399
Table 1 Some basic information of the six tested amino acids. Appellation
Abbreviation
Molecular weight (g/mol)
Molecular structure
O Glycine
Gly
H2N
75
CH C
OH
H O l-␣-Alanine
l-␣-Ala
H2N
89
CH C
OH
CH3
H dl-Alanine
dl-Ala
H3C
89
H COOH HOOC
C
C
CH3
NH2
NH2
O -Alanine
-Ala
Proline
Pro
89
H2N
H2C
H N
C
115
CH2 C
OH
O
OH
H2N Arginine
Arg
174
C NH
H N
NH2 H2 H2 H2 C C C CH C
O
OH
[22] and imidazolium [23] for the removal of SO2 have also been explored. However, commercial ILs are expensive and their preparation processes are complicated resulting in much higher costs and limiting their industrial applications. Recently, we demonstrated that l-␣-alanine supported on Al2 O3 could remove SO2 efficiently and effectively [24]. In this present paper, we have further expanded our work, and found that several water-soluble amino acid (Gly, l-␣-Ala, dl-Ala, -Ala, Pro and Arg) aqueous solutions showed excellent absorption performances on SO2 , such as high absorption capacity and good regeneration performance. So these amino acids could be alternative desulfurization absorbents. In fact, amino acids as biomolecules are the building blocks for peptides and proteins [25], and are usually used in biochemical and biological syntheses. To our knowledge, amino acids are seldom directly used for environment-related issues. As an extension herein, we have studied the desulfurization performance of six water-soluble amino acids (Gly, l-␣-Ala, dl-Ala, -Ala, Pro and Arg) aqueous solutions. Some basic information of these amino acids is presented in Table 1. The effects of the various factors on the desulfurization performance such as amino acid concentration, SO2 concentration, absorption and desorption temperature, and the initial pH value of absorbent are investigated by using -Ala solution. It is hoped that this work could enlighten researchers on the development of such desulfurization absorbents.
relatively long, for two specific advantages: (1) it prevents entrainment caused by the relatively high gas flow rate and (2) it facilitates refluxing and condensing during the desorption process (it requires as high as 150 ◦ C). Approximately 5.0 mL amino acid solution of appropriate concentration was placed in the tube reactor to remove SO2 from simulated flue gas. The simulated flue gas was a mixture of SO2 and N2 with appropriate concentration. The flow rate of the mixture gas was adjusted by a mass flow meter to 60 mL/min. The SO2 concentration changes at the outlet of the reactor were determined by iodine titration method (HJ/T 56-2000, a standard method of State Environmental Protection Administration of China). The amount of SO2 absorption capacity was expressed as milligram SO2 per milliliter amino acid solution (mg/mL). 2.2. SO2 desorption experiments After the saturation absorption (SO2 concentration at the outlet of the reactor was kept constant), the gas mixture was switched to N2 and the temperature increased to the required desorption temperature at a rate of 10 ◦ C/min; and the N2 flow was maintained until SO2 was no longer detected in the vent. Again, the concentration of SO2 at outlet of the reactor was determined by iodine titration method and the amount of desorbed SO2 expressed as before (mg/mL). Meanwhile, the absorbent was regenerated during the SO2 desorption process.
2. Materials and methods 2.3. 2.1. SO2 absorption experiments All the amino acids (biochemical reagent (BR)) used in the experiments were purchased from Sinopharm Chemical Reagent Co., Ltd, China. Experiments were performed in a vertically placed glass-made tube reactor (with 8 mm inside diameter and 39.5 cm long) under isothermal conditions. The tube reactor was made
13 C
NMR spectra
The original -Ala solution, -Ala solution saturated with SO2 , -Ala solution after SO2 desorption at 150 ◦ C and -Ala solution after ten continuous absorption–desorption cycles were mixed with heavy water (D2 O) in a volume ratio of 9:1 respectively. 13 C spectra were then recorded on a Varian Unity INOVA 100 MHz instrument.
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2.4. Ultraviolet spectra Ultraviolet (UV) spectra was use to monitor the original -Ala solution, the -Ala solution saturated with SO2 , -Ala solution after SO2 desorption at 150 ◦ C and the water saturated with SO2 . Samples of the solutions were diluted to the appropriate concentrations with deionized water and subsequently used for detection analyses. UV spectroscopy detection of the samples was carried out on TU-1900 (Purkinje General Instrument Co., Ltd., Beijing) spectrophotometers at a resolution of 1 nm. The spectra were recorded in the region of 190–400 nm. 3. Results and discussion 3.1. Activity of different amino acid solutions on SO2 removal Fig. 1a illustrates the activity of different amino acid (Gly, l-␣Ala, dl-Ala, -Ala, Pro and Arg) aqueous solutions on the removal of SO2 at 30 ◦ C. All the tested amino acid solutions exhibited good activity on the removal of SO2 compared with water. SO2 saturation uptake by -Ala solution was higher than that of the other ␣-amino acid solutions. This suggests that the structure of amino acids influences SO2 absorption to a great extent. In comparison to some other absorbents mentioned previously, for example, TMG ILs (SO2 saturation uptake is 305 mg/g, absorbed at 40 ◦ C, gas composition (volume fraction (vol.%): 8% SO2 and 92% N2 [21])) and imidazolium-based ILs (SO2 saturation uptake is 20.2 mg/g, absorbed at 20 ◦ C, gas composition: 10% SO2 and 90% N2 [23]), it is worth mentioning that SO2 saturation uptake by 1.0 M -Ala solution is relatively higher (about 390 mg/g). Fig. 1b shows that SO2 saturation uptake increases linearly with the increase in -Ala concentration. Increasing the concentration by 0.5 M increases the saturation uptake by an average of about 19.5 mg/mL. 3.2. Effects of SO2 concentration on SO2 removal by ˇ-Ala solution In Fig. 2, the effects of SO2 concentration on SO2 removal by -Ala solution at 30 ◦ C is shown. Fig. 2a displays the relationship between the rate of absorption and SO2 concentration. C0 represents the initial concentration of SO2 and Ct represents the concentration of SO2 at the exit of the reactor at time t. Ct /C0 could reflect the absorption rate in here. The saturation time (the time the absorption reaches saturation, i.e. at Ct /C0 = 1) decreased with the increase in the SO2 concentration. This could be attributed to the diffusion driving force increased with the increase in SO2 concentration, which increased SO2 absorption rate and thus
Fig. 1. (a) The activity of different amino acid aqueous solutions on SO2 removal. Absorption temperature: 30 ◦ C; concentration of amino acids: 1.0 M; gas composition (vol.%), SO2 : 4.1%; balance gas: N2 . (b) The activity of -Ala solution with different concentration on SO2 removal. Absorption temperature: 30 ◦ C; gas composition (vol.%): SO2 , 4.1%; balance gas: N2 .
promoted SO2 removal. Therefore, SO2 saturation uptake increased with the increasing of SO2 concentration as shown in Fig. 2b. However, a saturation uptake of about 26.8 mg/mL was realized when the concentration of SO2 was decreased to 0.8% revealing that the saturation uptake of SO2 decreased only by about 33% when the concentration of SO2 was decreased by about five times. This demonstrates that the -Ala solution could be employed to remove SO2 under relatively low SO2 concentrations.
Fig. 2. Effects of SO2 concentration on SO2 removal by -Ala solution. (a) Breakthrough curve, (b) variation of SO2 saturation uptake with SO2 concentration, absorption temperature: 30 ◦ C; concentration of -Ala (C-Ala ) is 1.0 M; balance gas: N2 .
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Fig. 3. Effects of temperature on SO2 removal by -Ala solution. C-Ala = 1.0 M; gas composition (vol.%): SO2 , 4.1%; balance gas: N2 .
3.3. Effects of absorption temperature on SO2 removal by ˇ-Ala solution Temperature is one of the main factors that affect the absorption of SO2 . Fig. 3 shows that SO2 saturation uptake was basically constant between 20 and 30 ◦ C but decreased (about 20.7%) with further increase in the absorption temperature to 60 ◦ C. This result clearly indicates that high absorption temperature was not favorable for SO2 removal, and as such room temperature range (20–30 ◦ C) was found to be the optimal condition for SO2 removal. This could be explained by the fact that high temperatures are not conducive for gaseous dissolution.
401
Fig. 5. Effects of initial pH value on SO2 removal by -Ala solution. Absorption temperature: 30 ◦ C; C-Ala = 1.0 M; gas composition (vol.%): SO2 , 4.1%; balance gas: N2 .
simply be due to SO2 physically dissolving in -Ala solution, so that it could be desorbed at the relatively low temperature of 30 ◦ C. Further still, SO2 desorption capacity and desorption rate are observed to respectively increased with increase in desorption temperature, and at a temperature of 150 ◦ C, SO2 had been desorbed almost completely. This result suggests that SO2 interacts with -Ala by different mechanisms besides by the simple physical dissolution. The SO2 dissolution by -Ala solution by different mechanisms requires different energy quantities for desorption and therefore, SO2 capacity increased with the increase in desorption temperature.
3.4. Effects of desorption temperature on SO2 desorption capacity
3.5. Effects of initial pH value on SO2 removal by ˇ-Ala solution
Fig. 4 reveals the relation between desorption temperature and SO2 desorption capacity. The SO2 desorption process was carried out after the SO2 absorption reached its saturation at 30 ◦ C. The processes were then operated at temperatures of 30, 60, 90, 120 and 150 ◦ C for 150 min under N2 flow. SO2 desorption capacity and desorption rate ((desorption capacity/saturation uptake) × 100%) were about 10.1 mg/mL and 25.2%, respectively, at the desorption temperature of 30 ◦ C. This implied that only a small portion of SO2 was desorbed at this temperature. The easily desorbed SO2 could
The initial pH value of -Ala solution was adjusted to 0.9, 2.9, 5.0 and 10.1, respectively, by addition of hydrochloric acid solution (1 M) or sodium hydroxide solution (1 M). The pH value of neutral -Ala solution (without any added acids or bases) was 6.8. Fig. 5 shows that SO2 saturation uptake is low when the adjusted pH of -Ala solution is low (4.5 mg/mL at 0.9 and 5.5 mg/mL at 2.9) due possibly, to the large H+ in the solution inhibiting the dissolution of SO2 . On the contrary, SO2 saturation uptake is higher with high adjusted pH (61.1 mg/mL at 10.1), even higher than that at the neutral pH (40.1 mg/mL at 6.8). However, desorption of SO2 was incomplete with desorption rate of 73.4% due partly to reaction SO2 with NaOH to form stable Na2 SO3 or NaHSO3 . Thus acidic -Ala solution does not promote SO2 absorption and alkaline Ala solution is not conducive for SO2 desorption. The neutral -Ala solution (pH = 6.8) is the optimal pH for SO2 removal. Amino acids (RCHNH2 COOH) in aqueous solution usually exist in the form of zwitterion (RCHNH3 + COO− ), and it may react with either H+ or OH− as illustrated in the example with -Ala reaction equation below: NH+ CHCH2 COO− + H+ NH+ CHCH2 COOH 3 3 CHCH2 COO− + OH− NH2 CHCH2 COO− + H2 O NH+ 3 Therefore, the absorption of SO2 in the neutral -Ala solution (pH = 6.8) may be described as below: 1. Dissolution of SO2 in water,
Fig. 4. Effects of desorption temperature on SO2 desorption capacity. Absorption temperature: 30 ◦ C, C-Ala = 1.0 M; gas composition (vol.%): SO2 , 4.1%; balance gas: N2 .
SO2 + H2 O H+ + HSO− 3
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Fig. 6. Regeneration performance of -Ala solution on SO2 removal. Absorption temperature: 30 ◦ C, C-Ala = 1.0 M; gas composition (vol.%): SO2 , 4.1%; balance gas: N2 .
2. Combination of carboxyl (from -Ala) with H+ , NH+ CH2 CH2 COO− 3
+
+H
NH+ CH2 CH2 COOH 3
This reaction greatly promotes the dissolution of SO2 . + 3. HSO− 3 further reacted with NH3 CH2 CH2 COOH, and this step will be discussed in detail in Section 3.7.
Fig. 7. 13 C NMR spectra (100 MHz) of original -Ala solution (1), -Ala solution saturation with SO2 (2), -Ala solution after SO2 desorption at 150 ◦ C (3) and Ala solution after ten continuous absorption–desorption cycles at the desorption temperature of 150 ◦ C (4). (Absorption temperature: 30 ◦ C; C-Ala = 1.0 M; gas composition (vol.%): SO2 , 4.1%; balance gas: N2 .)
solution was reversible and that -Ala was regenerated after SO2 desorption. This is consistent with the obtained experimental data.
3.6. Regeneration performance of ˇ-Ala solution The regeneration experiments were carried out at 30, 120 and 150 ◦ C desorption temperatures. Results of ten continuous absorption-desorption cycles of the -Ala solution are shown in Fig. 6, depicting constant SO2 desorption capacity at each temperature as a result of excellent regenerative ability. The indication is that the interaction between the SO2 and the -Ala solution is a reversible equilibrium reaction. 3.7. Absorption mechanism 3.7.1. 13 C NMR analysis The 13 C NMR analysis was used to explore the SO2 absorption mechanism. The 13 C NMR spectra of original -Ala solution (spectra 1), -Ala solution saturation with SO2 (spectra 2), Ala solution after SO2 desorption at 150 ◦ C (spectra 3) and -Ala solution after ten continuous absorption–desorption cycles at the desorption temperature of 150 ◦ C (spectra 4) are shown in Fig. 7. Three resonances observed at ı = 19.8, 22.9 and 164.5 ppm in original -Ala solution (NH3 + CH2 CH2 COO− ) are assigned to the carbon of CH2 , CH2 NH3 + and COO− respectively which shift to 17.5, 21.6 and 161.1 ppm respectively after saturation with SO2 . This chemical shift could be attributed to the result of SO2 interacting with -Ala. The results imply that there was no covalent bond formation between SO2 and -Ala, otherwise much bigger chemical shift would have been observed. The resonance of COO− changed by 3.4 ppm (from 164.5 to 161.1 ppm) after -Ala solution saturation with SO2 , while the resonances of CH2 and NH3 + CH2 changed by 2.3 and 1.3 ppm, respectively. This suggests that the SO2 interacted with COO− , and as a result, the chemical shift of COO− , CH2 and CH2 NH3 + showed a decreasing trend. All the three resonances were essentially restored to their original states after SO2 desorption at 150 ◦ C (spectra 3). Furthermore, the spectra of -Ala solution got restored after ten continuous absorption–desorption cycles (spectra 4). This result suggests that the absorption of SO2 by -Ala
3.7.2. UV spectra analysis Fig. 8 depicts the ultraviolet spectra of original -Ala solution (spectra 1), -Ala solution saturation with SO2 (spectra 2), -Ala solution after SO2 desorption at 150 ◦ C (spectra 3) and water saturation with SO2 (spectra 4). As mentioned earlier (Section 3.5), SO2 dissolves in water and ionized partly to form bisulfite ion. The two peaks observed at 276 nm and 204 nm in spectra 4 are assigned to SO2 and HSO3 − [26,27], respectively. The absorption peak of SO2 at 276 nm is also observed in the ultraviolet spectra of the -Ala solution saturated with SO2 (spectra 2) confirming that small portion of the SO2 indeed simply dissolves physically in the -Ala solution, as noted in Section 3.4. Comparatively, one notable difference between the ultraviolet spectra of water saturated with SO2 (spectra 4), and the spectra of -Ala solution saturated with
Fig. 8. Ultraviolet spectra of original -Ala solution (1), -Ala solution saturation with SO2 (2), -Ala solution after SO2 desorption at 150 ◦ C (3) and water saturation with SO2 (4). (Absorption temperature: 30 ◦ C, C-Ala = 1.0 M; gas composition (vol.%): SO2 : 4.1%; balance gas: N2 .).
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SO2 (spectra 2) was that the absorption peak of HSO3 − at 204 nm shifted to 210 nm. The absorption peak at 214 nm in the spectra of the original -Ala solution (spectra 1) is assigned to -Ala. This peak also shifted slightly after the -Ala solution is saturated with SO2 (spectra 2) because of the interaction of HSO3 − with -Ala. After the SO2 desorption at 150 ◦ C (spectra 3), the absorption peak (of SO2 ) at 276 nm disappeared and that of -Ala is restored (to 214 nm). Note that the -Ala solution saturated with SO2 was diluted 50 times before the UV scanning (otherwise its absorbance would have been beyond the instrument’s measurable range), while the rest of the other solutions were diluted 10 times. This suggests that the interaction between SO2 and -Ala has a hyperchromic effect on the UV absorbance of -Ala. Consequently, we speculate that SO2 dissolves in -Ala solution with a high saturation uptake because the interaction between them is not only through the simple physical dissolution; and that the interaction between them is not by covalent bonding but rather some weak forces of interactions, such as dispersion force, induction force, dipole–dipole force and hydrogen bonding. This observation is in agreement with that of van Dam who used organic solvents to remove SO2 [14]. In this study, we have shown that SO2 interacts with the carboxyl of -Ala to form hydrogen bond through a two-step process; the first step involves the dissolution of SO2 in water that is then ionized to generate hydrogen and bisulfite ions and the hydrogen ions combine with COO− to form COOH as mentioned earlier (Section 3.5). In the second step, hydrogen sulfite ions interact with COOH through hydrogen bonding. The process is summarized as:
In the above reaction with the dotted line representing hydrogen bond, the SO2 can be desorbed by heating. This partly explains why the SO2 saturation uptake by -Ala was higher than that of the other tested ␣-amino acids (Section 3.1). NH3 + is electron-withdrawing group, and has an electron withdrawing inductive effect on COOH. ␣-NH3 + is close to COOH and as such the inductive effect of electron withdrawing is much stronger in this case. This leads to the electrons being deflected from the COOH to NH3 + , which is not conducive for the formation of the hydrogen bond. On the contrary, -NH3 + is far from COOH, so it has little influence on the COOH group and therefore has little effect on the formation of hydrogen bond. In a nutshell, SO2 easily interacts with -Ala to form hydrogen bonds and therefore, the SO2 saturation uptake of -Ala is much higher than that observed for the other tested ␣-amino acids.
4. Conclusion Six water-soluble amino acids (Gly, l-␣-Ala, dl-Ala, -Ala, Pro and Arg) solutions were employed to remove SO2 from SO2 –N2 system. The effect of concentrations of the -Ala solution and SO2 , absorption and desorption temperatures and the initial pH of the -Ala solution on the desulfurization performance were investigated. Based on the results we conclude that: (1) The amino acid solutions have good activity on SO2 removal with -Ala solution exhibiting the highest saturation uptake. In fact, its saturation uptake increases with increasing concentrations of SO2 and -Ala solution. (2) While the optimal removal of SO2 is at room temperature (20–30 ◦ C), its optimal desorption is at high temperature (150 ◦ C). (3) Though acidic solution (pH = 0.9; 2.9) is not conducive for SO2 absorption, alkaline solution (pH = 10.1) is not conducive for SO2 desorption either; however, neutral solution (pH = 6.8) offers the optimal pH value for SO2 removal. (4) The interaction between SO2 and -Ala is not by covalent bonding but rather some weak interactive forces (such as dispersive, inductive, dipole-dipole
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and hydrogen bonding). (5) The absorbent (-Ala solution) posses excellent regenerative ability. Acknowledgements This project was supported by the Fujian Province Science & Technology Program of China (2010H6024). The authors are grateful for the State Key Laboratory for Physical Chemistry of Solid Surface of Xiamen University and Mr. Ibrahim Abdul-Rauf and Fazak Siyekwo for revision of the paper. References [1] M.J. Molina, L.T. Molina, Megacities and atmospheric pollution, J. Air Waste Manage. Assoc. 54 (2004) 644–680. [2] J.F. Izquierdo, C. Fite, F. Cunill, M. Iborra, J. Tejero, Kinetic study of the reaction between sulfur dioxide and calcium hydroxide at low temperature in a fixedbed reactor, J. Hazard. Mater. B76 (2000) 113–123. [3] K.T. Lee, A.R. Mohamed, S. Bhatia, K.H. Chu, Removal of sulfur dioxide by fly ash/CaO/CaSO4 sorbents, Chem. Eng. J. 114 (2005) 171–177. [4] I. Dahlan, K.T. Lee, A.H. Kamaruddin, A.R. Mohamed, Evaluation of various additives on the preparation of rice husk ash (RHA)/CaO-based sorbent for flue gas desulfurization (FGD) at low temperature, J. Hazard. Mater. 161 (2009) 570–574. [5] F.J. Gutiérrez Ortiz, A simple realistic modeling of full-scale wet limestone FGD units, Chem. Eng. J. 165 (2010) 426–439. [6] G.L. Hu, D.J. Kim, S. Wedel, J.P. Hansen, Review of the direct sulfation reaction of limestone, Prog. Energy Combust. 32 (2006) 386–407. [7] R.K. Srivastava, W. Jozewicz, C. Singer, SO2 scrubbing technologies: a review, Environ. Prog. 20 (2001) 219–228. [8] J. Huang, A. Riisager, P. Wasserscheid, R. Fehrmanna, Reversible physical absorption of SO2 by ionic liquids, Chem. Commun. 38 (2006) 4027–4029. [9] K. Stejskalova, I. Spirovová, E. Lippert, K. Mocek, Z. Bastl, A study of SO2 interaction with sodium carbonates by X-ray photoelectron spectroscopy, Appl. Surf. Sci. 103 (1996) 509–516. [10] C.F. Wu, S.J. Khang, T.C. Keener, S.K. Lee, A model for dry sodium bicarbonate duct injection flue gas desulfurization, Adv. Environ. Res. 8 (2004) 655–666. [11] S. Ebrahimi, C. Picioreanu, R. Kleerebezem, J.J. Heijnen, M.C.M. van Loosdrecht, Rate-based modelling of SO2 absorption into aqueous NaHCO3 /Na2 CO3 solutions accompanied by the desorption of CO2 , Chem. Eng. Sci. 58 (2003) 3589–3600. [12] G.H. Gleason, N.J. Montclair, A.C. Loonam, Recovery of sulphur dioxide, U.S. patent 21.106.453, 1940. [13] R.J. Demyanovich, S. Lynn, Prediction of infinite dilution activity coefficients of sulfur dioxide in organic solvents, J. Solution Chem. 20 (1991) 693–701. [14] M.H.H. van Dam, A.S. Lamine, D. Roizard, P. Lochon, C. Roizard, Selective sulfur dioxide removal using organic solvents, Ind. Eng. Chem. Res. 36 (1997) 4628–4637. [15] Z.G. Tang, C.C. Zhou, C. Chen, Studies on flue gas desulfurization by chemical absorption using an ethylenediamine–phosphoric acid solution, Ind. Eng. Chem. Res. 43 (2004) 6714–6722. [16] I.J. Uyanga, O.I. Raphael, Studies of SO2 and O2 induced degradation of aqueous MEA during CO2 capture from power plant flue gas streams, Ind. Eng. Chem. Res. 46 (2007) 2558–2566. [17] T. Supap, I. Raphael, P. Tontiwachwuthikul, C. Saiwan, Kinetics of sulfur dioxideand oxygen-induced degradation of aqueous monoethanolamine solution during CO2 absorption from power plant flue gas streams, Int. J. Green. Gas Con. 3 (2009) 133–142. [18] K. Fukumoto, M. Yoshizawa, H. Ohno, Room temperature ionic liquids from 20 natural amino acids, J. Am. Chem. Soc. 127 (2005) 2398–2399. [19] T. Welton, Room-temperature ionic liquids. Solvents for synthesis and catalysis, Chem. Rev. 99 (1999) 2071–2083. [20] Y.Y. Jiang, G.N. Wang, Z. Zhou, Y.T. Wu, J. Geng, Z.B. Zhang, Tetraalkylammonium amino acids as functionalized ionic liquids of low viscosity, Chem. Commun. 4 (2008) 505–507. [21] W.Z. Wu, B.X. Han, H.X. Gao, Z.M. Liu, T. Jiang, J. Huang, Desulfurization of flue gas: SO2 absorption by an ionic liquid, Angew. Chem. Int. Ed. 43 (2004) 2415–2417. [22] J. Huang, A. Riisager, R.W. Berg, R. Fehrmann, Tuning ionic liquids for high gas solubility and reversible gas sorption, J. Mol. Catal. A: Chem. 279 (2008) 170–176. [23] J. Huang, A. Riisager, P. Wasserscheid, R. Fehrmann, Reversible physical absorption of SO2 by ionic liquids, Chem. Commun. 38 (2006) 4027–4029. [24] R.P. Deng, L.S. Jia, Reversible removal of SO2 at low temperature by l-␣-alanine supported on ␥-Al2 O3 , Fuel 93 (2012) 385–390. [25] Z. Paˇı szti, L. Guczi, Amino acid adsorption on hydrophilic TiO2 : a sum frequency generation vibrational spectroscopy study, Vib. Spectrosc. 50 (2009) 48–56. [26] K. Fuwa, Bert L. Vallee, Molecular flame absorption spectrometry for sulfur, Anal. Chem. 41 (1969) 188–190. [27] A. Syty, Determination of sulfur dioxide by ultraviolet absorption spectrometry, Anal. Chem. 45 (1973) 1744–1747.