- Email: [email protected]
Mg−Al layered double hydroxide intercalated with CO32– and its recyclability for treatment of SO2
Applied Clay Science 183 (2019) 105349
Contents lists available at ScienceDirect
Applied Clay Science journal homepage: www.elsevier.com/locate/clay
Research Paper
Mg−Al layered double hydroxide intercalated with CO32– and its recyclability for treatment of SO2
T
Tomohito Kameda , Masahito Tochinai, Shogo Kumagai, Toshiaki Yoshioka ⁎
Graduate School of Environmental Studies, Tohoku University, 6−6−07 Aoba, Aramaki, Aoba−ku, Sendai 980-8579, Japan
ARTICLE INFO
ABSTRACT
Keywords: SO2 CO3·Mg − Al layered double hydroxide Recycling
Mg − Al layered double hydroxide intercalated with CO32– (CO3·Mg − Al LDH) was successfully used to remove SO2 from a gas mixture at 140, 170, and 200 °C. The removal of SO2 with CO3·Mg − Al LDH occurred via surface adsorption of SO2 and anion exchange between CO32– in the interlayer of Mg − Al LDH and SO2 from the gas mixture. The removal of SO2 followed the order of Ca(OH)2 > CO3·Mg − Al LDH (Mg/Al = 2) > CO3·Mg − Al LDH (Mg/Al = 4). The removal of SO2 also followed the order of CO3·Mg − Al LDH > Mg (OH)2 > SO4·Mg − Al LDH. CO3·Mg − Al LDH was regenerated from SO4·Mg-Al LDH using a Na2CO3 solution through anion exchange between SO42− and CO32−, leading to its recyclability for removal of SO2 from a gas mixture three times, demonstrating the feasibility of recycling the material.
1. Introduction In Japan, approximately 80% of general waste is incinerated. General waste includes S components, which upon incineration produce acid gases such as SOx. The lime–gypsum method is well known as a flue gas desulfurization method for SOx (Chu and Rochelle, 1989; Sakai et al., 2002). Although this effective treatment of SOx is currently essential to reduce air pollution, it poses other environmental problems such as increased amounts of desulfurized gypsum and reduced landfill lifetimes. Therefore, a new treatment process is required. Mg − Al layered double hydroxide (Mg − Al LDH) is an anion exchanger that can intercalate various types of anions in its interlayer space (Ingram and Taylor, 1967; Allmann, 1968; Miyata, 1983; Cavani et al., 1991; Mills et al., 2012). The general chemical formula is 3+ n− [Mg2+ )x/n·mH2O, where x denotes the Al/(Mg + Al) 1−xAlx (OH)2](A molar ratio (0.20 ≦ x ≦ 0.33) and An− is a counterion. Mg − Al LDH with intercalated CO32– (CO3·Mg − Al LDH) can be transformed to Mg − Al oxide by calcination at ~500 °C, according to the following reaction:
Mg1 x Alx (OH)2 (CO3 ) x/2
Mg1 x Alx O1 + x/2 + x/2CO2 + H2 O
The resulting Mg − Al oxide can be rehydrated and combined with anions to reconstruct the LDH structure according to the following reaction: Mg1 x Alx O1 + x/2 + x/n An + (1 + x/2)H2 O
⁎
Mg1 x Alx (OH) 2 Ax/n + xOH
In our previous study, a unique treatment method for SO2 was developed using MgeAl oxide in a wet process (Kameda et al., 2012). In that study, a Mg − Al oxide slurry was prepared by mixing Mg − Al oxide with H2O. Dissolution of SO2 into the slurry produced SO3·Mg − Al LDH, which was then calcined at 900–1000 °C to yield H2SO4, MgO, and MgAl2O4. The decomposition products, MgO and MgAl2O4, are useful refractory materials that can be recovered from the process, thereby reducing the amount of waste. Although this wet process is effective, the preparation of Mg − Al oxide is cumbersome. In this study, the effectiveness of using CO3·Mg − Al LDH in a dry process to treat SO2 is examined. Certain parameters were analyzed experimentally to better understand their effect on SO2 treatment efficiency. These parameters include the quantity of CO3·Mg − Al LDH used in the process, temperature, and SO2 concentration in the waste stream, which were analyzed to prepare a breakthrough curve. The chemicals used for acid gas treatment in incineration facilities are not recycled, and this is a factor that also makes disposal methods very stringent. We therefore considered SO2 treatment using CO3·Mg − Al LDH, as shown in Fig. 1. Furthermore, SO2 removal using CO3·Mg − Al LDH was compared to SO4·Mg − Al LDH, Mg(OH)2, and Ca(OH)2. Assuming that the reaction between SO2 and CO3·Mg-Al LDH yields SO3·Mg-Al LDH and SO4·Mg-Al LDH, we propose that regeneration of CO3·Mg − Al LDH via anion exchange is possible. We therefore synthesized SO4·Mg-Al LDH by coprecipitation, and studied the influence of the stoichiometric quantity of CO32– on anion exchange with SO42− on MgeAl LDH. SO42− desorption was analyzed kinetically and at
Corresponding author. E-mail address: [email protected] (T. Kameda).
https://doi.org/10.1016/j.clay.2019.105349 Received 25 March 2019; Received in revised form 23 October 2019; Accepted 24 October 2019 0169-1317/ © 2019 Elsevier B.V. All rights reserved.
Applied Clay Science 183 (2019) 105349
T. Kameda, et al.
solution preheated to a predetermined temperature for experiments to probe the effect of CO32−, and to analyze kinetics and the equilibrium state. In order to examine the influence of the stoichiometric quantity of CO32−, 0.2 g of SO4·Mg-Al LDH was added to the Na2CO3 solution at 0.5 to 2.5 times the stoichiometric amount. For kinetic analysis, SO4·Mg-Al LDH with a stoichiometric quantity of 1.0 (0.30 g) was added to 0.025 M Na2CO3 solution. For equilibrium analysis, 0.2 g of SO4·Mg-Al LDH was added to 0.005–0.5 M Na2CO3 solution. The Erlenmeyer flask was placed in a shaker set to 10–60 °C and shaken for 0.5–180 min. After measuring the pH of the post-reaction solution, the solid and liquid were separated using a 0.45 μm membrane filter. A Dionex DX-120 ion chromatograph and a Dionex model AS-12A column (eluent: 2.7 mM sodium carbonate and 0.3 mM sodium bicarbonate, flow rate: 1.3 mL/min) was used to measure the SO42− concentration in the filtrate, the total organic carbon (TOC; multi N/C 2110S) was used to determine the C concentration, and inductively coupled plasma atomic emission spectroscopy (ICP-AES; iCAP6500Duo) was used to measure the elution of Mg2+ and Al3+ monitored at 279.5 and 308.2 nm, respectively. After drying the products under reduced pressure at 40 °C, phase identification was carried out by XRD.
SO2
CO3•Mg−Al LDH
SO3•Mg−Al LDH SO4•Mg−Al LDH Anion exchange CO32−
SO32−, SO42−
Fig. 1. Proposed treatment of SO2 using CO3 •Mg−Al LDH.
equilibrium. As shown in Fig. 1, CO3·Mg-Al LDH was used to treat SO2, then we performed anion exchange using a Na2CO3 solution to regenerate CO3·Mg-Al LDH, and then the treatment of the acid gas mixture was repeated in order to assess its recyclability. 2. Experimental
2.3. Recyclability
2.1. Treatment of SO2
Firstly, SO2 treatment with CO3·Mg − Al LDH was conducted according to Section 2–1. For this, 50 ppm SO2 was treated with 10.0 times the stoichiometric quantity of CO3·Mg − Al LDH (Mg/Al = 2 or 4) according to Eq. (1) at 170 °C for 90 min. Next, regeneration of CO3·Mg − Al LDH after SO2 treatment was conducted according to Section 2–2. A Na2CO3 solution (20 mL) with twice the stoichiometric quantity according to Eq. (2) was added to CO3·Mg − Al LDH (Mg/ Al = 2 or 4) after SO2 treatment, and was shaken at 30 °C for 180 min. This procedure was repeated 3 times.
CO3·Mg − Al LDH (Mg/Al = 2) and CO3·Mg − Al LDH (Mg/Al = 4) were prepared by coprecipitation, according to our previously reported procedure (Kameda et al., 2016). The chemical composition and properties of the LDH are displayed in Table S1. The removal of SO2 by CO3·Mg − Al LDH (Mg/Al = 2) and CO3·Mg − Al LDH (Mg/Al = 4) can be expressed with Eq. (1). Mg1 x Alx (OH) 2 (CO3 ) x/2 ·mH2 O + x/2SO2
Mg1
x Alx (OH)2
(SO3 ) x/2 + x/2CO2 + mH2 O (x = 0.20, 0.33)
(1)
3. Results and discussion
The amount of CO3·Mg − Al LDH required to treat gaseous SO2 was 5.0–15.0 times the stoichiometric quantities indicated in Eq. (1). In another experiment, the amount of CO3·Mg − Al LDH required to treat gaseous SO2 was determined to be 0.25–0.50 g. A quartz tube reactor (inner diameter: 16 mm; length: 450 mm) was prepared with a glass wool insert placed in the midsection of the tube. The glass wool insert was included to enable adjustment of CO3·Mg − Al LDH in the middle of the hot reactor zone, additionally ensuring passage of the SO2 gas stream. An electric furnace was set at 140–200 °C, and SO2 gas was added to a flow of N2 gas to obtain SO2 gas concentrations of 50–150 ppm at a linear velocity of 1.0 m/min for 540 min. After stopping the SO2 gas mixture flow, the quartz tube reactor was purged with N2. To evaluate the effectiveness of SO2 removal using CO3·Mg − Al LDH, the same experiment was performed using SO4·Mg − Al LDH, Mg (OH)2, and Ca(OH)2. Mg(OH)2 and Ca(OH)2 are of any trademark. SO2 in the evolved gas was quantified using a gas analyzer (testo 350XL), and the products were identified by X-ray diffraction (XRD) analyses using a Rigaku RINT 2200 diffractometer with Cu Kα radiation at 40 kV and 20 mA (scanning rate of 2°/min).
3.1. Treatment of SO2 Fig. 2 shows the effect of the quantity of CO3·Mg − Al LDH on the removal of SO2. Both CO3·Mg − Al LDH with Mg/Al molar ratios of 2 and 4 were found to remove SO2 from the gas stream. The removal of SO2 increased with increasing amount of CO3·Mg − Al LDH. SO2 removal at 15.0 times the stoichiometric quantity was found to be 61.8 and 58.2% for Mg/Al = 2 and Mg/Al = 4, respectively. Fig. S1 shows
100
SO2 removal [%]
80
2.2. Anion desorption
Mg/Al=4.0 60
40
20
SO4·Mg-Al LDH was synthesized by coprecipitation with a Mg/Al mole ratio of 2:1 and a SO42− content of 16.5 wt%. Eq. (2) shows the theoretical reaction for anion exchange between SO4·Mg-Al LDH and CO32−. Mg 0.67 Al 0.33 (OH) 2 (SO4 ) 0.17 + 0.17CO32
Mg/Al=2.0
0 5.0
10.0
15.0
CO3•Mg−Al LDH quantity [stoichiometric quantity]
Mg 0.67Al 0.33 (OH)2 (CO3 )0.17 + 0.17SO4 2
(2)
Fig. 2. Effect of CO3•Mg-Al LDH quantity on the removal of SO2. SO2concentration:50[ppm];Temperature:170[ºC];Linear velocity:1.0[m/min];Time: 90[min]
Briefly, in a 50 mL screw-cap Erlenmeyer flask, a predetermined amount of SO4·Mg-Al LDH was added to 20 mL of an aqueous Na2CO3 2
Applied Clay Science 183 (2019) 105349
T. Kameda, et al.
100
100
Mg/Al=2.0
Mg/Al=2.0 80
80
Mg/Al=4.0
SO2 removal [%]
SO2 removal [%]
Mg/Al=4.0 60
40
20
60
40
20
0
0
140
170 Temperarture [ºC]
200
50
Fig. 3. Effect of temperature on the removal of SO2. SO2concentration:50[ppm];CO3•Mg-Al LDH quantity[10.0-times the stoichiometric quantity];Linear velocity:1.0[m/min];Time:90[min].
100 150 SO2 concentration [ppm]
Fig. 4. Effect of SO2 concentration on the removal of SO2. Temperature:170[ºC];CO3•Mg-Al LDH quantity[10.0-times the stoichiometric quantity];Linear velocity:1.0[m/min];Time:90[min]
the XRD patterns of the products after SO2 removal from a gas stream with three different stoichiometric quantities (5.0, 10.0, and 15.0) of CO3·Mg − Al LDH (Mg/Al = 2) and CO3·Mg − Al LDH (Mg/Al = 4). The products are identified as hydrotalcite (JCPDS card 22–700), a naturally occurring hydroxycarbonate of magnesium and aluminum (Mg6Al2(OH)16CO3·4H2O) with a LDH structure. This indicates that the structure of CO3·Mg − Al LDH was maintained during SO2 removal. In both Mg/Al molar ratios, the d003 values, which indicate the interlayer spacing of Mg − Al LDH, decreased with increasing stoichiometric quantity. This suggests the desorption of CO32– in the interlayer of Mg − Al LDH and the intercalation of SO32− according to Eq. (1). Fig. 3 shows the effect of temperature on the removal of SO2. For both CO3·Mg − Al LDH, the removal of SO2 increased slightly with increasing temperature. SO2 removal at 200 °C was found to be 52.0 and 70.0% for Mg/Al = 2 and 4, respectively. In an actual incineration plant, the lime–gypsum method is performed around 170 °C. Therefore, the following experiments were performed at 170 °C. Fig. 4 shows the effect of SO2 concentration on the removal of SO2. For both CO3·Mg − Al LDH, SO2 removal increased with increasing SO2 concentration. This can be attributed to the increase in contact between CO3·Mg − Al LDH and SO2. To evaluate the relative effectiveness of CO3·Mg − Al LDH, its removal efficiency was compared to that of Ca (OH)2. Fig. 5 shows the variation in the removal of SO2 by CO3·Mg − Al LDH (Mg/Al = 2), CO3·Mg − Al LDH (Mg/Al = 4), and Ca(OH)2 as a function of time, using 0.25 g of CO3·Mg − Al LDH and Ca(OH)2. The SO2 removal by Ca(OH)2 was 100% at 150 min and 94.5% at 180 min. The SO2 removal by CO3·Mg − Al LDH (Mg/Al = 2) was 100% at 30 min and 89.1% at 60 min. The SO2 removal by CO3·Mg − Al LDH (Mg/Al = 4) was 87.3% at 30 min. The removal of SO2 was found to be high (> 85%) for all three substances in the order of Ca (OH)2 > CO3·Mg − Al LDH (Mg/Al = 2) > CO3·Mg − Al LDH (Mg/ Al = 4). Fig. S2 shows the variation in the removal of SO2 by 0.50 g of the two CO3·Mg − Al LDH (Mg/Al = 2 and 4). The removal of SO2 increased with increasing amounts of CO3·Mg − Al LDH; the SO2 removal with CO3·Mg − Al LDH (Mg/Al = 2) was 100% at 60 min and 96.4% at 90 min, and that with CO3·Mg − Al LDH (Mg/Al = 4) was 100% at 90 min and 92.7% at 120 min. Fig. 6 shows the variation in residual SO2 concentration after treatment with CO3·Mg − Al LDH (Mg/Al = 2), SO4·Mg − Al LDH (Mg/ Al = 2), and Mg(OH)2 as a function of time. For all materials, the residual SO2 concentrations were low during the initial reaction stage, indicating that all three were able to remove SO2 from the gas. SO2 removal by Mg(OH)2 is attributed to surface adsorption of SO2.
100
SO2 removal [%]
80
60 Mg/Al=2.0 40
Mg/Al=4.0 Ca(OH)2 Ca(OH)2
20
0 0
100
200
300 400 Time [min]
500
600
Fig. 5. Variation in the removal of SO2 by CO3•Mg-Al LDH(Mg/Al=2), CO3•Mg-Al LDH(Mg/Al=4) and Ca(OH)2 with time. SO2 concentration:50[ppm];Tem perature:170[ºC];Linear velocity:1.0[m/min];CO3•Mg-Al LDH and Ca(OH)2 quantity:0.25[g]
Residual SO2 concentration [ppm]
60 50 SO4•Mg−Al LDH 40 Mg(OH)2
30 20
CO3•Mg−Al LDH 10 0 0
20
40
60
Time [min] Fig. 6. Variation in the residual SO2 concentration after the treatment by CO3•Mg−Al LDH (Mg/Al = 2), SO4•Mg−Al LDH (Mg/Al = 2) and Mg(OH)2 with time. SO2 concentration: 50 [ppm]; Temperature: 170 [ºC]; Linear velocity: 1.0 [m/min]; Quantity: 15.0-times the stoichiometric quantity for CO3•Mg−Al LDH quantity
3
Applied Clay Science 183 (2019) 105349
T. Kameda, et al.
Likewise, SO2 removal by SO4·Mg − Al LDH is also attributed to surface adsorption of SO2 because the anion exchange in Eq. (1) did not occur. The residual SO2 concentrations after treatment with CO3·Mg − Al LDH, Mg(OH)2, and SO4·Mg − Al LDH for 60 min were found to be 23, 32, and 47 ppm, respectively. SO2 removal was found to be high for all three substances in the order of CO3·Mg − Al LDH > Mg (OH)2 > SO4·Mg − Al LDH. The structure of CO3·Mg − Al LDH is similar to those of SO4·Mg − Al LDH and Mg(OH)2. Therefore, the surface adsorption of SO2 is considered to contribute to SO2 removal by CO3·Mg − Al LDH. In addition, the difference between SO2 removal by CO3·Mg − Al LDH and Mg(OH)2 or SO4·Mg − Al LDH is attributed to the anion exchange process shown in Eq. (1). In summary, the removal of SO2 by CO3·Mg − Al LDH is due to both surface adsorption of SO2 and anion exchange, as shown in Eq. (1). Fig. 8. Change over time in SO42- desorption rate at each temperature. Na2CO3 concentration: 0.025 [mol/L]; CO32- stoichiometric quantity: 1.0.
3.2. Anion desorption 3.2.1. Influence of the stoichiometry of CO32– on SO42− desorption Figs. 7, S3, and S4 show the influence of the stoichiometric quantity of CO32– on SO42− desorption, pH, and Mg elution. Even upon adding 0.5-times the stoichiometric quantity of CO32−, 58.2% SO42− desorption was confirmed. Here, OH– is also thought to have intercalated in the same way. This is supported by the decrease in pH from the initial pH, suggesting adsorption of OH−. Furthermore, as the stoichiometric quantity of CO32– increased, SO42− desorption increased, and at 2.5 times the stoichiometric quantity, 93.4% desorption of SO42− was possible. Furthermore, decreasing the pH from the initial pH made it possible to reduce Mg elution. Fig. S5 shows the XRD patterns of the products. While maintaining a hydrotalcite structure, the basic plane spacing of d003 = 1.05 nm in the original SO4·Mg-Al LDH decreased as the stoichiometric quantity of CO32– increased, and when the stoichiometric quantity of CO32– was 1.0 or more, the d003 matched with that of CO3·Mg-Al LDH. In the case of 0.5 times the stoichiometric quantity used (Fig. S5b), both CO32– and SO42− were thought to be intercalated in the interlayer of MgeAl LDH.
constant [min−1]. When Eq. (3) is integrated, the SO42− desorption can be represented by x to obtain Eq. (4).
ln(1
dqt dt
= k (qe
qt )2
(5)
Here, k is the apparent reaction constant [(min・mmol・g)−1], and qe and qt [mmol/g] are the amount desorbed at equilibrium and at time t [min], respectively. By integrating Eq. (5) and rearranging it, we obtain Eq. (6).
t 1 1 = + t qt kqe 2 qe
(6)
Fig. S8 shows the plot of reaction time versus t/qt for the first stage of the reaction shown in Fig. 8. The correlation coefficient suggests a good linear relationship between the reaction time and t/qt, indicating that the desorption of SO42− from SO4·Mg-Al LDH (Mg/Al = 2:1) due to CO32– exhibits pseudo-second order behavior. 3.2.3. Adsorption of CO32– and HCO3−/desorption isotherm of SO42− Figs. S9 shows the adsorption/desorption isotherm of SO42− desorption and adsorption of CO32– and HCO3– during the reaction of SO4·Mg-Al LDH and CO32−. We fitted the data using the Langmuir equation and the Freundlich equation. The Langmuir equation is expressed as
(3)
Here, t is the reaction time ([min] and k is the apparent rate 100
SO42- desorption [%]
(4)
Application of the pseudo-first order rate equation at the initial stage of the reaction (based on Fig. 8) results in a plot of −ln(1 − x) as a function of the reaction time (Fig. S7). Based on the correlation coefficient, we found that a good linear relationship cannot be established between the reaction time and − ln(1 − x), and by having an intercept, it did not follow pseudo-first order reaction kinetics. Next, we applied the following pseudo-second order rate equation:
3.2.2. Kinetic analysis Figs. 8 and S6 show the change in SO42− desorption over time and pH at each temperature. At 10 and 30 °C, no significant differences are observed in SO42− desorption; however, at 60 °C, the SO42− desorption is rapid. Furthermore, the pH continued to decrease up to the highest temperature. From these results, we analyzed the rate of the desorption reaction. The pseudo-first order equation is given by
d [Cl ] = k [Cl ] dt
x ) = kt
80
qe =
60
KL qm Ce 1 + KL Ce
(7)
Here, qe is the amount of adsorbed SO2 at equilibrium [mmol/g], Ce is the equilibrium concentration [mmol/L] of carbon C (CO32– and HCO3−), qm is the maximum adsorption and desorption [mmol/g], and KL is the adsorption/desorption equilibrium constant [L/mmol]. Rearranging Eq. (7) gives Eq. (8).
40 20
Ce 1 1 = + Ce qe KL qm qm
0 0.5
1
1.5 2 2.5 CO32- stoichiometric quantity
(8)
Fig. S10 shows a plot of Ce/qe vs. Ce. Both the adsorption and desorption processes show a good linear relationship between Ce/qe and Ce and fit the Langmuir equation. The maximum adsorption (qm) and adsorption equilibrium constant (KL) determined from the slope and
Fig. 7. Influence of CO32- stoichiometric quantity on the SO42- desorption. Temperature: 30 [ºC]; Time: 60 [min] 4
Applied Clay Science 183 (2019) 105349
T. Kameda, et al.
intercept of the line are 1.18 mmol/g and 1.86 L/mmol, respectively. Furthermore, the maximum desorption qm and equilibrium constant KL are 1.54 mmol/L and 1.68 L/mmol, respectively. Here, based on (maximum adsorption) – (maximum desorption), we can state that 0.36 mmol/L of OH– has adsorbed. Here, a relationship can be established between the equilibrium adsorption, the adsorption equilibrium constant KL, and the standard Gibbs free energy ΔG0 (kJ/mol), as shown in Eq. (9).
Mg/Al=2.0 80 SO2 removal [%]
G0 =
100
(9)
RTlnKL
The standard Gibbs free energy ΔG0 at 30 °C through the adsorption and desorption pathways are both −18.7 kJ/mol. The negative ΔG0 value indicates that the adsorption process of CO32– and HCO3−, as well as the desorption process of SO42− that occurs simultaneously, are spontaneous reactions. Next, the Freundlich equation is an empirical equation determined through experiments, and is shown in Eq. (10). 1
1 log Ce n
1
standard entropy ΔS0 are calculated to be −6.58 kJ/mol and 3.94 × 10−2 kJ/(mol·K), respectively. Since the standard enthalpy ΔH0 is negative, SO42− desorption is exothermic.
(11)
3.3. Recyclability Fig. 10 shows the influence of the number of recycling cycles on SO2 removal. SO2 removal in the first treatment using CO3·Mg − Al LDH with Mg/Al = 2 and 4 are 54.6 and 51.8%, respectively. However, by the third treatment, this decreased marginally to 44.7 and 36.2%, respectively, thus demonstrating the effectiveness of CO3·Mg-Al LDH recyclability for SO2 treatment when used three times.
3.2.4. Thermodynamic analysis Figs. 9 and S12 show the desorption isotherm, and Ce/qe vs. Ce plots for SO42− desorption at 10, 30, and 60 °C. At each temperature, Ce/qe and Ce show a good linear relationship. The desorption equilibrium constants KL determined from the slopes of the lines are 1.86, 1.70, and 1.23 L/mmol at 10, 30, and 60 °C, respectively. Furthermore, the standard Gibbs free energies ΔG0 at 10, 30, and 60 °C are −17.7, −18.7, and − 19.7 kJ/mol, respectively. Here, the relationship shown in Eq. (12) is established between the desorption equilibrium constant KL, standard enthalpy ΔH0, and standard entropy ΔS0.
RTlnKL =
H0
4. Conclusions CO3·Mg − Al LDH with Mg/Al molar ratios of 2 and 4 were used to remove SO2 from a gas mixture. The removal of SO2 followed the order of Ca(OH)2 > CO3·Mg − Al LDH (Mg/Al = 2) > CO3·Mg − Al LDH (Mg/Al = 4). The removal of SO2 also followed the order of CO3·Mg − Al LDH > Mg(OH)2 > SO4·Mg − Al LDH, suggesting that removal of SO2 by CO3·Mg − Al LDH is due to both surface adsorption of SO2 and anion exchange (shown in Eq. (1)). The desorption of SO42− from SO4·Mg − Al LDH was achieved through anion exchange of CO32– with SO42− from MgeAl LDH. CO3·Mg−Al LDH can be repeatedly used in the removal of SO2 from a gas, which suggests its feasibility as a recyclable material for SO2 removal.
(12)
T S0
Further rearrangement results in Eq. (13).
ln KL =
S0 R
H0 RT
2 3 Number of recycling cycles
Fig. 10. Influence of the number of recycling cycles on the SO2 removal. SO2 concentration: 50 [ppm]; CO3•Mg−Al LDH quantity [10.0-times the stoichiometric quantity]; Temperature: 170 [ºC].
Fig. S11 shows a plot of logqe vs. logCe. A good correlation is not established between logCe and logqe, which implies that an approximation using the Freundlich equation cannot be applied.
G0 =
40
0
Here, KF is the adsorption and desorption equilibrium constant [L/ mol] and n is a constant. Taking the common logarithm of both sides of Eq. (10), we obtain Eq. (11).
log q e = log KF +
60
20
(10)
q e = KF Ce n
Mg/Al=4.0
(13) −1
Fig. S13 shows a plot of lnKL vs. T . A good linear relationship exists between lnKL and T−1, and the standard enthalpy ΔH0 and
Declaration of Competing Interest The authors declare that they have no known competing financial interests or personal relationships that could have appeared to influence the work reported in this paper. Acknowledgments This research was supported by the Environment Research and Technology Development Fund (3K163007) of the Ministry of the Environment, Japan. Appendix A. Supplementary data Fig. 9. Desorption isotherms on the SO42- desorption. Na2CO3 concentration: 0.005~0.5 [mol/L]; Time: 180 [min]
Supplementary data to this article can be found online at https:// doi.org/10.1016/j.clay.2019.105349. 5
Applied Clay Science 183 (2019) 105349
T. Kameda, et al.
References
Chemosphere 88, 250–254. Kameda, T., Tochinai, M., Yoshioka, T., 2016. Treatment of hydrochloric acid using Mg-Al layered double hydroxide intercalated with carbonate. J. Ind. Eng. Chem. 39, 21–26. Mills, S.J., Christy, A.G., Genin, J.M.R., Kameda, T., Colombo, F., 2012. Nomenclature of the hydrotalcite supergroup: Natural layered double hydroxides. Mineral. Mag. 76, 1289–1336. Miyata, S., 1983. Anion-exchange properties of hydrotalcite-like compounds. Clay Clay Miner. 31, 305–311. Sakai, M., Su, C., Sasaoka, E., 2002. Simultaneous removal of SOx and NOx using slaked lime at low temperature. Ind. Eng. Chem. Res. 41, 5029–5033.
Allmann, R., 1968. The crystal structure of pyroaurite. Acta Crystallogr. B24, 972–977. Cavani, F., Trifiro, F., Vaccari, A., 1991. Hydrotalcite-type anionic clays: Preparation, properties, and applications. Catal. Today 11, 173–301. Chu, P., Rochelle, G.T., 1989. Removal of SO2 and NOX from stack gas by reaction with calcium hydroxide solids. J. Air Pollut. Control Assoc. 39, 175–179. Ingram, L., Taylor, H.F.W., 1967. The crystal structures of sjogrenite and pyroaurite. Mineral. Mag. 36, 465–479. Kameda, T., Kodama, A., Fubasami, Y., Kumagai, S., Yoshioka, T., 2012. Removal of SO2 with a Mg-Al oxide slurry via reconstruction of a Mg-Al layered double hydroxide.
6
Contents lists available at ScienceDirect
Applied Clay Science journal homepage: www.elsevier.com/locate/clay
Research Paper
Mg−Al layered double hydroxide intercalated with CO32– and its recyclability for treatment of SO2
T
Tomohito Kameda , Masahito Tochinai, Shogo Kumagai, Toshiaki Yoshioka ⁎
Graduate School of Environmental Studies, Tohoku University, 6−6−07 Aoba, Aramaki, Aoba−ku, Sendai 980-8579, Japan
ARTICLE INFO
ABSTRACT
Keywords: SO2 CO3·Mg − Al layered double hydroxide Recycling
Mg − Al layered double hydroxide intercalated with CO32– (CO3·Mg − Al LDH) was successfully used to remove SO2 from a gas mixture at 140, 170, and 200 °C. The removal of SO2 with CO3·Mg − Al LDH occurred via surface adsorption of SO2 and anion exchange between CO32– in the interlayer of Mg − Al LDH and SO2 from the gas mixture. The removal of SO2 followed the order of Ca(OH)2 > CO3·Mg − Al LDH (Mg/Al = 2) > CO3·Mg − Al LDH (Mg/Al = 4). The removal of SO2 also followed the order of CO3·Mg − Al LDH > Mg (OH)2 > SO4·Mg − Al LDH. CO3·Mg − Al LDH was regenerated from SO4·Mg-Al LDH using a Na2CO3 solution through anion exchange between SO42− and CO32−, leading to its recyclability for removal of SO2 from a gas mixture three times, demonstrating the feasibility of recycling the material.
1. Introduction In Japan, approximately 80% of general waste is incinerated. General waste includes S components, which upon incineration produce acid gases such as SOx. The lime–gypsum method is well known as a flue gas desulfurization method for SOx (Chu and Rochelle, 1989; Sakai et al., 2002). Although this effective treatment of SOx is currently essential to reduce air pollution, it poses other environmental problems such as increased amounts of desulfurized gypsum and reduced landfill lifetimes. Therefore, a new treatment process is required. Mg − Al layered double hydroxide (Mg − Al LDH) is an anion exchanger that can intercalate various types of anions in its interlayer space (Ingram and Taylor, 1967; Allmann, 1968; Miyata, 1983; Cavani et al., 1991; Mills et al., 2012). The general chemical formula is 3+ n− [Mg2+ )x/n·mH2O, where x denotes the Al/(Mg + Al) 1−xAlx (OH)2](A molar ratio (0.20 ≦ x ≦ 0.33) and An− is a counterion. Mg − Al LDH with intercalated CO32– (CO3·Mg − Al LDH) can be transformed to Mg − Al oxide by calcination at ~500 °C, according to the following reaction:
Mg1 x Alx (OH)2 (CO3 ) x/2
Mg1 x Alx O1 + x/2 + x/2CO2 + H2 O
The resulting Mg − Al oxide can be rehydrated and combined with anions to reconstruct the LDH structure according to the following reaction: Mg1 x Alx O1 + x/2 + x/n An + (1 + x/2)H2 O
⁎
Mg1 x Alx (OH) 2 Ax/n + xOH
In our previous study, a unique treatment method for SO2 was developed using MgeAl oxide in a wet process (Kameda et al., 2012). In that study, a Mg − Al oxide slurry was prepared by mixing Mg − Al oxide with H2O. Dissolution of SO2 into the slurry produced SO3·Mg − Al LDH, which was then calcined at 900–1000 °C to yield H2SO4, MgO, and MgAl2O4. The decomposition products, MgO and MgAl2O4, are useful refractory materials that can be recovered from the process, thereby reducing the amount of waste. Although this wet process is effective, the preparation of Mg − Al oxide is cumbersome. In this study, the effectiveness of using CO3·Mg − Al LDH in a dry process to treat SO2 is examined. Certain parameters were analyzed experimentally to better understand their effect on SO2 treatment efficiency. These parameters include the quantity of CO3·Mg − Al LDH used in the process, temperature, and SO2 concentration in the waste stream, which were analyzed to prepare a breakthrough curve. The chemicals used for acid gas treatment in incineration facilities are not recycled, and this is a factor that also makes disposal methods very stringent. We therefore considered SO2 treatment using CO3·Mg − Al LDH, as shown in Fig. 1. Furthermore, SO2 removal using CO3·Mg − Al LDH was compared to SO4·Mg − Al LDH, Mg(OH)2, and Ca(OH)2. Assuming that the reaction between SO2 and CO3·Mg-Al LDH yields SO3·Mg-Al LDH and SO4·Mg-Al LDH, we propose that regeneration of CO3·Mg − Al LDH via anion exchange is possible. We therefore synthesized SO4·Mg-Al LDH by coprecipitation, and studied the influence of the stoichiometric quantity of CO32– on anion exchange with SO42− on MgeAl LDH. SO42− desorption was analyzed kinetically and at
Corresponding author. E-mail address: [email protected] (T. Kameda).
https://doi.org/10.1016/j.clay.2019.105349 Received 25 March 2019; Received in revised form 23 October 2019; Accepted 24 October 2019 0169-1317/ © 2019 Elsevier B.V. All rights reserved.
Applied Clay Science 183 (2019) 105349
T. Kameda, et al.
solution preheated to a predetermined temperature for experiments to probe the effect of CO32−, and to analyze kinetics and the equilibrium state. In order to examine the influence of the stoichiometric quantity of CO32−, 0.2 g of SO4·Mg-Al LDH was added to the Na2CO3 solution at 0.5 to 2.5 times the stoichiometric amount. For kinetic analysis, SO4·Mg-Al LDH with a stoichiometric quantity of 1.0 (0.30 g) was added to 0.025 M Na2CO3 solution. For equilibrium analysis, 0.2 g of SO4·Mg-Al LDH was added to 0.005–0.5 M Na2CO3 solution. The Erlenmeyer flask was placed in a shaker set to 10–60 °C and shaken for 0.5–180 min. After measuring the pH of the post-reaction solution, the solid and liquid were separated using a 0.45 μm membrane filter. A Dionex DX-120 ion chromatograph and a Dionex model AS-12A column (eluent: 2.7 mM sodium carbonate and 0.3 mM sodium bicarbonate, flow rate: 1.3 mL/min) was used to measure the SO42− concentration in the filtrate, the total organic carbon (TOC; multi N/C 2110S) was used to determine the C concentration, and inductively coupled plasma atomic emission spectroscopy (ICP-AES; iCAP6500Duo) was used to measure the elution of Mg2+ and Al3+ monitored at 279.5 and 308.2 nm, respectively. After drying the products under reduced pressure at 40 °C, phase identification was carried out by XRD.
SO2
CO3•Mg−Al LDH
SO3•Mg−Al LDH SO4•Mg−Al LDH Anion exchange CO32−
SO32−, SO42−
Fig. 1. Proposed treatment of SO2 using CO3 •Mg−Al LDH.
equilibrium. As shown in Fig. 1, CO3·Mg-Al LDH was used to treat SO2, then we performed anion exchange using a Na2CO3 solution to regenerate CO3·Mg-Al LDH, and then the treatment of the acid gas mixture was repeated in order to assess its recyclability. 2. Experimental
2.3. Recyclability
2.1. Treatment of SO2
Firstly, SO2 treatment with CO3·Mg − Al LDH was conducted according to Section 2–1. For this, 50 ppm SO2 was treated with 10.0 times the stoichiometric quantity of CO3·Mg − Al LDH (Mg/Al = 2 or 4) according to Eq. (1) at 170 °C for 90 min. Next, regeneration of CO3·Mg − Al LDH after SO2 treatment was conducted according to Section 2–2. A Na2CO3 solution (20 mL) with twice the stoichiometric quantity according to Eq. (2) was added to CO3·Mg − Al LDH (Mg/ Al = 2 or 4) after SO2 treatment, and was shaken at 30 °C for 180 min. This procedure was repeated 3 times.
CO3·Mg − Al LDH (Mg/Al = 2) and CO3·Mg − Al LDH (Mg/Al = 4) were prepared by coprecipitation, according to our previously reported procedure (Kameda et al., 2016). The chemical composition and properties of the LDH are displayed in Table S1. The removal of SO2 by CO3·Mg − Al LDH (Mg/Al = 2) and CO3·Mg − Al LDH (Mg/Al = 4) can be expressed with Eq. (1). Mg1 x Alx (OH) 2 (CO3 ) x/2 ·mH2 O + x/2SO2
Mg1
x Alx (OH)2
(SO3 ) x/2 + x/2CO2 + mH2 O (x = 0.20, 0.33)
(1)
3. Results and discussion
The amount of CO3·Mg − Al LDH required to treat gaseous SO2 was 5.0–15.0 times the stoichiometric quantities indicated in Eq. (1). In another experiment, the amount of CO3·Mg − Al LDH required to treat gaseous SO2 was determined to be 0.25–0.50 g. A quartz tube reactor (inner diameter: 16 mm; length: 450 mm) was prepared with a glass wool insert placed in the midsection of the tube. The glass wool insert was included to enable adjustment of CO3·Mg − Al LDH in the middle of the hot reactor zone, additionally ensuring passage of the SO2 gas stream. An electric furnace was set at 140–200 °C, and SO2 gas was added to a flow of N2 gas to obtain SO2 gas concentrations of 50–150 ppm at a linear velocity of 1.0 m/min for 540 min. After stopping the SO2 gas mixture flow, the quartz tube reactor was purged with N2. To evaluate the effectiveness of SO2 removal using CO3·Mg − Al LDH, the same experiment was performed using SO4·Mg − Al LDH, Mg (OH)2, and Ca(OH)2. Mg(OH)2 and Ca(OH)2 are of any trademark. SO2 in the evolved gas was quantified using a gas analyzer (testo 350XL), and the products were identified by X-ray diffraction (XRD) analyses using a Rigaku RINT 2200 diffractometer with Cu Kα radiation at 40 kV and 20 mA (scanning rate of 2°/min).
3.1. Treatment of SO2 Fig. 2 shows the effect of the quantity of CO3·Mg − Al LDH on the removal of SO2. Both CO3·Mg − Al LDH with Mg/Al molar ratios of 2 and 4 were found to remove SO2 from the gas stream. The removal of SO2 increased with increasing amount of CO3·Mg − Al LDH. SO2 removal at 15.0 times the stoichiometric quantity was found to be 61.8 and 58.2% for Mg/Al = 2 and Mg/Al = 4, respectively. Fig. S1 shows
100
SO2 removal [%]
80
2.2. Anion desorption
Mg/Al=4.0 60
40
20
SO4·Mg-Al LDH was synthesized by coprecipitation with a Mg/Al mole ratio of 2:1 and a SO42− content of 16.5 wt%. Eq. (2) shows the theoretical reaction for anion exchange between SO4·Mg-Al LDH and CO32−. Mg 0.67 Al 0.33 (OH) 2 (SO4 ) 0.17 + 0.17CO32
Mg/Al=2.0
0 5.0
10.0
15.0
CO3•Mg−Al LDH quantity [stoichiometric quantity]
Mg 0.67Al 0.33 (OH)2 (CO3 )0.17 + 0.17SO4 2
(2)
Fig. 2. Effect of CO3•Mg-Al LDH quantity on the removal of SO2. SO2concentration:50[ppm];Temperature:170[ºC];Linear velocity:1.0[m/min];Time: 90[min]
Briefly, in a 50 mL screw-cap Erlenmeyer flask, a predetermined amount of SO4·Mg-Al LDH was added to 20 mL of an aqueous Na2CO3 2
Applied Clay Science 183 (2019) 105349
T. Kameda, et al.
100
100
Mg/Al=2.0
Mg/Al=2.0 80
80
Mg/Al=4.0
SO2 removal [%]
SO2 removal [%]
Mg/Al=4.0 60
40
20
60
40
20
0
0
140
170 Temperarture [ºC]
200
50
Fig. 3. Effect of temperature on the removal of SO2. SO2concentration:50[ppm];CO3•Mg-Al LDH quantity[10.0-times the stoichiometric quantity];Linear velocity:1.0[m/min];Time:90[min].
100 150 SO2 concentration [ppm]
Fig. 4. Effect of SO2 concentration on the removal of SO2. Temperature:170[ºC];CO3•Mg-Al LDH quantity[10.0-times the stoichiometric quantity];Linear velocity:1.0[m/min];Time:90[min]
the XRD patterns of the products after SO2 removal from a gas stream with three different stoichiometric quantities (5.0, 10.0, and 15.0) of CO3·Mg − Al LDH (Mg/Al = 2) and CO3·Mg − Al LDH (Mg/Al = 4). The products are identified as hydrotalcite (JCPDS card 22–700), a naturally occurring hydroxycarbonate of magnesium and aluminum (Mg6Al2(OH)16CO3·4H2O) with a LDH structure. This indicates that the structure of CO3·Mg − Al LDH was maintained during SO2 removal. In both Mg/Al molar ratios, the d003 values, which indicate the interlayer spacing of Mg − Al LDH, decreased with increasing stoichiometric quantity. This suggests the desorption of CO32– in the interlayer of Mg − Al LDH and the intercalation of SO32− according to Eq. (1). Fig. 3 shows the effect of temperature on the removal of SO2. For both CO3·Mg − Al LDH, the removal of SO2 increased slightly with increasing temperature. SO2 removal at 200 °C was found to be 52.0 and 70.0% for Mg/Al = 2 and 4, respectively. In an actual incineration plant, the lime–gypsum method is performed around 170 °C. Therefore, the following experiments were performed at 170 °C. Fig. 4 shows the effect of SO2 concentration on the removal of SO2. For both CO3·Mg − Al LDH, SO2 removal increased with increasing SO2 concentration. This can be attributed to the increase in contact between CO3·Mg − Al LDH and SO2. To evaluate the relative effectiveness of CO3·Mg − Al LDH, its removal efficiency was compared to that of Ca (OH)2. Fig. 5 shows the variation in the removal of SO2 by CO3·Mg − Al LDH (Mg/Al = 2), CO3·Mg − Al LDH (Mg/Al = 4), and Ca(OH)2 as a function of time, using 0.25 g of CO3·Mg − Al LDH and Ca(OH)2. The SO2 removal by Ca(OH)2 was 100% at 150 min and 94.5% at 180 min. The SO2 removal by CO3·Mg − Al LDH (Mg/Al = 2) was 100% at 30 min and 89.1% at 60 min. The SO2 removal by CO3·Mg − Al LDH (Mg/Al = 4) was 87.3% at 30 min. The removal of SO2 was found to be high (> 85%) for all three substances in the order of Ca (OH)2 > CO3·Mg − Al LDH (Mg/Al = 2) > CO3·Mg − Al LDH (Mg/ Al = 4). Fig. S2 shows the variation in the removal of SO2 by 0.50 g of the two CO3·Mg − Al LDH (Mg/Al = 2 and 4). The removal of SO2 increased with increasing amounts of CO3·Mg − Al LDH; the SO2 removal with CO3·Mg − Al LDH (Mg/Al = 2) was 100% at 60 min and 96.4% at 90 min, and that with CO3·Mg − Al LDH (Mg/Al = 4) was 100% at 90 min and 92.7% at 120 min. Fig. 6 shows the variation in residual SO2 concentration after treatment with CO3·Mg − Al LDH (Mg/Al = 2), SO4·Mg − Al LDH (Mg/ Al = 2), and Mg(OH)2 as a function of time. For all materials, the residual SO2 concentrations were low during the initial reaction stage, indicating that all three were able to remove SO2 from the gas. SO2 removal by Mg(OH)2 is attributed to surface adsorption of SO2.
100
SO2 removal [%]
80
60 Mg/Al=2.0 40
Mg/Al=4.0 Ca(OH)2 Ca(OH)2
20
0 0
100
200
300 400 Time [min]
500
600
Fig. 5. Variation in the removal of SO2 by CO3•Mg-Al LDH(Mg/Al=2), CO3•Mg-Al LDH(Mg/Al=4) and Ca(OH)2 with time. SO2 concentration:50[ppm];Tem perature:170[ºC];Linear velocity:1.0[m/min];CO3•Mg-Al LDH and Ca(OH)2 quantity:0.25[g]
Residual SO2 concentration [ppm]
60 50 SO4•Mg−Al LDH 40 Mg(OH)2
30 20
CO3•Mg−Al LDH 10 0 0
20
40
60
Time [min] Fig. 6. Variation in the residual SO2 concentration after the treatment by CO3•Mg−Al LDH (Mg/Al = 2), SO4•Mg−Al LDH (Mg/Al = 2) and Mg(OH)2 with time. SO2 concentration: 50 [ppm]; Temperature: 170 [ºC]; Linear velocity: 1.0 [m/min]; Quantity: 15.0-times the stoichiometric quantity for CO3•Mg−Al LDH quantity
3
Applied Clay Science 183 (2019) 105349
T. Kameda, et al.
Likewise, SO2 removal by SO4·Mg − Al LDH is also attributed to surface adsorption of SO2 because the anion exchange in Eq. (1) did not occur. The residual SO2 concentrations after treatment with CO3·Mg − Al LDH, Mg(OH)2, and SO4·Mg − Al LDH for 60 min were found to be 23, 32, and 47 ppm, respectively. SO2 removal was found to be high for all three substances in the order of CO3·Mg − Al LDH > Mg (OH)2 > SO4·Mg − Al LDH. The structure of CO3·Mg − Al LDH is similar to those of SO4·Mg − Al LDH and Mg(OH)2. Therefore, the surface adsorption of SO2 is considered to contribute to SO2 removal by CO3·Mg − Al LDH. In addition, the difference between SO2 removal by CO3·Mg − Al LDH and Mg(OH)2 or SO4·Mg − Al LDH is attributed to the anion exchange process shown in Eq. (1). In summary, the removal of SO2 by CO3·Mg − Al LDH is due to both surface adsorption of SO2 and anion exchange, as shown in Eq. (1). Fig. 8. Change over time in SO42- desorption rate at each temperature. Na2CO3 concentration: 0.025 [mol/L]; CO32- stoichiometric quantity: 1.0.
3.2. Anion desorption 3.2.1. Influence of the stoichiometry of CO32– on SO42− desorption Figs. 7, S3, and S4 show the influence of the stoichiometric quantity of CO32– on SO42− desorption, pH, and Mg elution. Even upon adding 0.5-times the stoichiometric quantity of CO32−, 58.2% SO42− desorption was confirmed. Here, OH– is also thought to have intercalated in the same way. This is supported by the decrease in pH from the initial pH, suggesting adsorption of OH−. Furthermore, as the stoichiometric quantity of CO32– increased, SO42− desorption increased, and at 2.5 times the stoichiometric quantity, 93.4% desorption of SO42− was possible. Furthermore, decreasing the pH from the initial pH made it possible to reduce Mg elution. Fig. S5 shows the XRD patterns of the products. While maintaining a hydrotalcite structure, the basic plane spacing of d003 = 1.05 nm in the original SO4·Mg-Al LDH decreased as the stoichiometric quantity of CO32– increased, and when the stoichiometric quantity of CO32– was 1.0 or more, the d003 matched with that of CO3·Mg-Al LDH. In the case of 0.5 times the stoichiometric quantity used (Fig. S5b), both CO32– and SO42− were thought to be intercalated in the interlayer of MgeAl LDH.
constant [min−1]. When Eq. (3) is integrated, the SO42− desorption can be represented by x to obtain Eq. (4).
ln(1
dqt dt
= k (qe
qt )2
(5)
Here, k is the apparent reaction constant [(min・mmol・g)−1], and qe and qt [mmol/g] are the amount desorbed at equilibrium and at time t [min], respectively. By integrating Eq. (5) and rearranging it, we obtain Eq. (6).
t 1 1 = + t qt kqe 2 qe
(6)
Fig. S8 shows the plot of reaction time versus t/qt for the first stage of the reaction shown in Fig. 8. The correlation coefficient suggests a good linear relationship between the reaction time and t/qt, indicating that the desorption of SO42− from SO4·Mg-Al LDH (Mg/Al = 2:1) due to CO32– exhibits pseudo-second order behavior. 3.2.3. Adsorption of CO32– and HCO3−/desorption isotherm of SO42− Figs. S9 shows the adsorption/desorption isotherm of SO42− desorption and adsorption of CO32– and HCO3– during the reaction of SO4·Mg-Al LDH and CO32−. We fitted the data using the Langmuir equation and the Freundlich equation. The Langmuir equation is expressed as
(3)
Here, t is the reaction time ([min] and k is the apparent rate 100
SO42- desorption [%]
(4)
Application of the pseudo-first order rate equation at the initial stage of the reaction (based on Fig. 8) results in a plot of −ln(1 − x) as a function of the reaction time (Fig. S7). Based on the correlation coefficient, we found that a good linear relationship cannot be established between the reaction time and − ln(1 − x), and by having an intercept, it did not follow pseudo-first order reaction kinetics. Next, we applied the following pseudo-second order rate equation:
3.2.2. Kinetic analysis Figs. 8 and S6 show the change in SO42− desorption over time and pH at each temperature. At 10 and 30 °C, no significant differences are observed in SO42− desorption; however, at 60 °C, the SO42− desorption is rapid. Furthermore, the pH continued to decrease up to the highest temperature. From these results, we analyzed the rate of the desorption reaction. The pseudo-first order equation is given by
d [Cl ] = k [Cl ] dt
x ) = kt
80
qe =
60
KL qm Ce 1 + KL Ce
(7)
Here, qe is the amount of adsorbed SO2 at equilibrium [mmol/g], Ce is the equilibrium concentration [mmol/L] of carbon C (CO32– and HCO3−), qm is the maximum adsorption and desorption [mmol/g], and KL is the adsorption/desorption equilibrium constant [L/mmol]. Rearranging Eq. (7) gives Eq. (8).
40 20
Ce 1 1 = + Ce qe KL qm qm
0 0.5
1
1.5 2 2.5 CO32- stoichiometric quantity
(8)
Fig. S10 shows a plot of Ce/qe vs. Ce. Both the adsorption and desorption processes show a good linear relationship between Ce/qe and Ce and fit the Langmuir equation. The maximum adsorption (qm) and adsorption equilibrium constant (KL) determined from the slope and
Fig. 7. Influence of CO32- stoichiometric quantity on the SO42- desorption. Temperature: 30 [ºC]; Time: 60 [min] 4
Applied Clay Science 183 (2019) 105349
T. Kameda, et al.
intercept of the line are 1.18 mmol/g and 1.86 L/mmol, respectively. Furthermore, the maximum desorption qm and equilibrium constant KL are 1.54 mmol/L and 1.68 L/mmol, respectively. Here, based on (maximum adsorption) – (maximum desorption), we can state that 0.36 mmol/L of OH– has adsorbed. Here, a relationship can be established between the equilibrium adsorption, the adsorption equilibrium constant KL, and the standard Gibbs free energy ΔG0 (kJ/mol), as shown in Eq. (9).
Mg/Al=2.0 80 SO2 removal [%]
G0 =
100
(9)
RTlnKL
The standard Gibbs free energy ΔG0 at 30 °C through the adsorption and desorption pathways are both −18.7 kJ/mol. The negative ΔG0 value indicates that the adsorption process of CO32– and HCO3−, as well as the desorption process of SO42− that occurs simultaneously, are spontaneous reactions. Next, the Freundlich equation is an empirical equation determined through experiments, and is shown in Eq. (10). 1
1 log Ce n
1
standard entropy ΔS0 are calculated to be −6.58 kJ/mol and 3.94 × 10−2 kJ/(mol·K), respectively. Since the standard enthalpy ΔH0 is negative, SO42− desorption is exothermic.
(11)
3.3. Recyclability Fig. 10 shows the influence of the number of recycling cycles on SO2 removal. SO2 removal in the first treatment using CO3·Mg − Al LDH with Mg/Al = 2 and 4 are 54.6 and 51.8%, respectively. However, by the third treatment, this decreased marginally to 44.7 and 36.2%, respectively, thus demonstrating the effectiveness of CO3·Mg-Al LDH recyclability for SO2 treatment when used three times.
3.2.4. Thermodynamic analysis Figs. 9 and S12 show the desorption isotherm, and Ce/qe vs. Ce plots for SO42− desorption at 10, 30, and 60 °C. At each temperature, Ce/qe and Ce show a good linear relationship. The desorption equilibrium constants KL determined from the slopes of the lines are 1.86, 1.70, and 1.23 L/mmol at 10, 30, and 60 °C, respectively. Furthermore, the standard Gibbs free energies ΔG0 at 10, 30, and 60 °C are −17.7, −18.7, and − 19.7 kJ/mol, respectively. Here, the relationship shown in Eq. (12) is established between the desorption equilibrium constant KL, standard enthalpy ΔH0, and standard entropy ΔS0.
RTlnKL =
H0
4. Conclusions CO3·Mg − Al LDH with Mg/Al molar ratios of 2 and 4 were used to remove SO2 from a gas mixture. The removal of SO2 followed the order of Ca(OH)2 > CO3·Mg − Al LDH (Mg/Al = 2) > CO3·Mg − Al LDH (Mg/Al = 4). The removal of SO2 also followed the order of CO3·Mg − Al LDH > Mg(OH)2 > SO4·Mg − Al LDH, suggesting that removal of SO2 by CO3·Mg − Al LDH is due to both surface adsorption of SO2 and anion exchange (shown in Eq. (1)). The desorption of SO42− from SO4·Mg − Al LDH was achieved through anion exchange of CO32– with SO42− from MgeAl LDH. CO3·Mg−Al LDH can be repeatedly used in the removal of SO2 from a gas, which suggests its feasibility as a recyclable material for SO2 removal.
(12)
T S0
Further rearrangement results in Eq. (13).
ln KL =
S0 R
H0 RT
2 3 Number of recycling cycles
Fig. 10. Influence of the number of recycling cycles on the SO2 removal. SO2 concentration: 50 [ppm]; CO3•Mg−Al LDH quantity [10.0-times the stoichiometric quantity]; Temperature: 170 [ºC].
Fig. S11 shows a plot of logqe vs. logCe. A good correlation is not established between logCe and logqe, which implies that an approximation using the Freundlich equation cannot be applied.
G0 =
40
0
Here, KF is the adsorption and desorption equilibrium constant [L/ mol] and n is a constant. Taking the common logarithm of both sides of Eq. (10), we obtain Eq. (11).
log q e = log KF +
60
20
(10)
q e = KF Ce n
Mg/Al=4.0
(13) −1
Fig. S13 shows a plot of lnKL vs. T . A good linear relationship exists between lnKL and T−1, and the standard enthalpy ΔH0 and
Declaration of Competing Interest The authors declare that they have no known competing financial interests or personal relationships that could have appeared to influence the work reported in this paper. Acknowledgments This research was supported by the Environment Research and Technology Development Fund (3K163007) of the Ministry of the Environment, Japan. Appendix A. Supplementary data Fig. 9. Desorption isotherms on the SO42- desorption. Na2CO3 concentration: 0.005~0.5 [mol/L]; Time: 180 [min]
Supplementary data to this article can be found online at https:// doi.org/10.1016/j.clay.2019.105349. 5
Applied Clay Science 183 (2019) 105349
T. Kameda, et al.
References
Chemosphere 88, 250–254. Kameda, T., Tochinai, M., Yoshioka, T., 2016. Treatment of hydrochloric acid using Mg-Al layered double hydroxide intercalated with carbonate. J. Ind. Eng. Chem. 39, 21–26. Mills, S.J., Christy, A.G., Genin, J.M.R., Kameda, T., Colombo, F., 2012. Nomenclature of the hydrotalcite supergroup: Natural layered double hydroxides. Mineral. Mag. 76, 1289–1336. Miyata, S., 1983. Anion-exchange properties of hydrotalcite-like compounds. Clay Clay Miner. 31, 305–311. Sakai, M., Su, C., Sasaoka, E., 2002. Simultaneous removal of SOx and NOx using slaked lime at low temperature. Ind. Eng. Chem. Res. 41, 5029–5033.
Allmann, R., 1968. The crystal structure of pyroaurite. Acta Crystallogr. B24, 972–977. Cavani, F., Trifiro, F., Vaccari, A., 1991. Hydrotalcite-type anionic clays: Preparation, properties, and applications. Catal. Today 11, 173–301. Chu, P., Rochelle, G.T., 1989. Removal of SO2 and NOX from stack gas by reaction with calcium hydroxide solids. J. Air Pollut. Control Assoc. 39, 175–179. Ingram, L., Taylor, H.F.W., 1967. The crystal structures of sjogrenite and pyroaurite. Mineral. Mag. 36, 465–479. Kameda, T., Kodama, A., Fubasami, Y., Kumagai, S., Yoshioka, T., 2012. Removal of SO2 with a Mg-Al oxide slurry via reconstruction of a Mg-Al layered double hydroxide.
6