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A study of CO2 precipitation method considering an ionic CO2 and Ca(OH)2 slurry
Energy 75 (2014) 624e629
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A study of CO2 precipitation method considering an ionic CO2 and Ca(OH)2 slurry Sangwon Park, Hoyong Jo, Dongwoo Kang, Jinwon Park* Department of Chemical & Biomolecular Engineering, Yonsei University, 50 Yonsei-ro, Seodaemun-gu, Seoul 120-749, Republic of Korea
a r t i c l e i n f o
a b s t r a c t
Article history: Received 18 January 2014 Received in revised form 2 July 2014 Accepted 9 August 2014 Available online 28 August 2014
CCS (carbon capture and storage) is the most popular technology used for the reduction of CO2 in the post-combustion stage. However, the CCS process has some disadvantages including uncertainty about the stability of the land that is used to store the separated CO2. Consequently, CCU (carbon capture and utilization) technologies have recently received increased attention as a possible replacement for CCS. In this study, we utilized CO2 fixation methods by using the metal carbonate mechanism. We selected 5 and 30 wt% MEA (mono-ethanolamine) solutions to rapidly make a carbonate and Ca(OH)2 slurry. In all of the experiments, normal temperature and pressure conditions were maintained (except during desorption to check for residual CO2 in the MEA solution). Consequently, most of the CO2 was converted to carbonate. The MEA converted CO2 to ionic CO2 and rapidly created calcium carbonate. Also the formed solids that were observed were determined to be CaCO3 and Ca(OH)2 by X-ray diffractometry. Also, the MEA solution could be reused to absorb CO2. Therefore, we have confirmed the development of our suggested CCS process. This process has the ability not only to reuse emitted CO2, but it can also be employed to reuse construction wastes that include heavy metals. © 2014 Elsevier Ltd. All rights reserved.
Keywords: CCS (carbon capture and storage) process CO2 mineralization PCC (precipitate calcium carbonate)
1. Introduction It is widely accepted that GHGs (greenhouse gases) are responsible for the greenhouse effect [1]. Most GHGs are emitted by human activities such as from coal-fired power plants that use fossil fuels [2]. Thus, many researchers have been exploring various CO2 reduction methods to curtail the greenhouse effect. The CCS (carbon capture and storage) process is a well-known technology that is widely employed to reduce CO2 emissions after combustion [3,4]. CCS appears particularly well-suited for the relatively large, high concentration CO2 emissions found in many types of industry [4]. For example, CCS technology research has focused on coal-fired power plants in Korea. However, the CCS process has some disadvantages. First, a large area is required to separate CO2 from the absorbents. As previously stated, it is beneficial to apply CCS technologies to industries that produce high CO2 emissions because the amount of CO2 that is separated is equal to the amount that is absorbed using this process. According to the KEPRI (Korea Electric Power Research Institute), a 10 MW coal-fired power plant emits 80,000 tons of CO2 in a single year [5]. A large area of land would be
* Corresponding author. Tel.: þ82 2 364 1807; fax: þ82 2 312 6401. E-mail address: [email protected] (J. Park). http://dx.doi.org/10.1016/j.energy.2014.08.036 0360-5442/© 2014 Elsevier Ltd. All rights reserved.
required to store this CO2. This poses a problem for countries such as Korea, which have a small total land area. Secondly, there are issues regarding the stability of storing CO2 in the land or ocean [6,7]. Separated CO2 is stored underground or in the ocean using absorbents. However, there is the possibility that the stored CO2 could be reemitted into the atmosphere after physical changes to the storage area such as damage caused from earthquakes or volcanoes. Finally, there are also economic problems. The CCS process is made up of two parts: absorption and desorption [8]. Between these, CO2 desorption uses the majority of the energy (more than 70%) [6]. In the desorption process, CO2 in the absorbents is separated using heat or pressure. As a result, the CCS process consumes a large amount of energy [9]. Therefore, some researchers have attempted to find ways to reduce energy consumption during desorption [10,11]. We believe that if certain improvements are made (e.g. researching alternative absorbents and additives to reduce energy consumption), CCS can become superior to other CO2 reduction methods from an economic standpoint [12]. CCU (carbon capture and utilization) technologies have been studied since the 1990s as a potential method to overcome the aforementioned shortcomings of CCS [13]. The goal of CCU is to reuse captured CO2 via biological, physical, and/or chemical means. In the biological case, captured CO2 can be converted into renewable energy or acetic acid using algae [6,10]. In the physical and
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chemical cases, CO2 fixation and reuse methods were used. These methods focused on the chemical conversion mechanism. For example, CO2 gas was converted to a solid (e.g. calcium, barium, or magnesium carbonate) by combining it with metal ions [14,15]. Solid CO2 conversion byproducts have better chemical stability compared to gases or ionic liquids. Following these methods, the emitted CO2 can be stored underground or in the ocean and ultimately reused [6]. With regard to the metal carbonate mechanism, CO2 fixation methods were classified as using either high- or lowenergy depending on the total energy consumption. In general, high-energy consumption methods were used to directly emit CO2 and some of the metals [16,17]. These methods require high temperature and pressure to make the metal carbonate because metal ions, including metal materials, need a certain amount of activation energy to react with CO2 gas. According to Lackner et al., calcium and magnesium carbonates from serpentine require temperatures above 500 C and a pressure greater than 340 bars [18]. In addition, magnesium carbonate from magnesium hydroxide requires 540 C and 45 bars [19]. These methods appear to have a significant disadvantage from the standpoint of energy consumption. However, some industries, including power plants and the cement industry, do have sufficient energy to support high-temperature fixation methods. In addition, the waste heat generated in this process could be used [10]. Conversely, low-energy CO2 fixation does not require high temperature or high pressure [6,10]. The methods mentioned above use conversion materials. In other words, emitted CO2 is converted to ionic CO2 by conversion materials such as amine, ammonia, and alkaline solutions, which are well-known CO2 absorbents [9e11]. Acid gas CO2 is converted to ionic CO2 during absorption in an alkaline solution. An alkaline solution generally creates ionic CO2 via CO2 absorption, carbonate, bicarbonate, and carbamate, respectively. Then, the ionic CO2 reacts with metal ions. Therefore, ionic CO2 could form a metal carbonate more easily than gaseous CO2 because it has a lower activation energy. According to Sangwon Park et al., low-energy consumption CO2 fixation methods are superior to the CCS process because they maintain normal conditions of 30 C and 1 bar. They also do not require a desorption process [6,10]. In addition, most of the CO2 absorbed in their amine solution was converted to metal carbonates. This means that additional energy was not required in desorption and regeneration of the amine solution. Therefore, we believe that this method is preferable from an economic standpoint. However, when considering the aforementioned studies, we believe that two important points should be emphasized. First, the rate at which CO2 is converted to ionic CO2 and the absorbed CO2 conversion kinetics must be taken into account. Second, it must be determined whether or not the fixation sources can provide a sufficient amount of ionic CO2 to produce a metal carbonate. A limited supply of metal ions might not produce enough metal carbonate. In other words, converted CO2 would remain in solution because the total CO2 loading is larger than the amount of supplied metal ions. In this regard, this method is similar to CCS. Additionally, a low CO2 conversion rate could incur a larger reactor size when enough metal is provided. This means that additional energy consumption would be required (similar to the high-energy consumption process for rapidly making metal carbonates in smaller reactors). Thus, we have taken these facts into consideration for CO2 fixation under low-energy consumption conditions. We explored ways to make a carbonate using CCU under normal temperatures and pressures. We selected an amine solution in order to convert CO2, and we selected Ca(OH)2 as the metal source. We chose the amine solution because it has a high CO2 conversion rate. In addition, we assumed that this study could improve CCS by considering desorption; thus, we determined that an amine solution would make a suitable conversion material. Finally, we
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selected MEA (mono-ethanolamine) because it was widely used in most CCS processes and has been verified by several researchers [3e5,20]. Ca(OH)2 was easily obtained from concrete and construction wastes. Therefore, this study verified the possibility of simultaneously obtaining CO2 and Ca(OH)2 emitted from construction waste under normal conditions. 2. Experimental The experimental apparatus is shown in Fig. 1. We removed the remaining CO2 by using high-purity N2 gas (99.999%). The highpurity N2 gas caught the residual CO2 in the experiment, and the mixing gas flowed through a Teflon tube. The mixed gases were detected with a CO2 analyzer at 30 s intervals. We maintained this process for 30 min. After that, we assumed that the CO2 concentration was 15 V/V% in a balance of N2. N2 and CO2 gases were controlled by MFC (mass flow rate), and the flow rates were 1870 and 300 ml/min, respectively. Also, the purity of the CO2 and N2 gases used in the absorption-desorption process were both 99.9%. Each gas was mixed by a saturated reactor which flowed into the conversion reactor. The conversion reactor was filled with 5 and 30 wt% MEA solutions. Each MEA solution was 400 ml and balanced on water using the corresponding weight ratio. MEA was purchased from SigmaeAldrich and was more than 99% pure. Simulated gas continuously flowed into the CO2 analyzer, and its concentration was analyzed at 30 s intervals. After the conversion reaction, the simulated gas was flowed through a cold condenser (15 C). This was used to maintain the concentration of the conversion solution via trapping any water loss. A reaction temperature of 30 C was maintained in the saturated and conversion reactors by a thermo water bath. Finally, we assumed that the conversion experiment was finished when the outlet CO2 concentration reached 15 V/V%. After the conversion experiment was completed, the CO2 precipitation experiment began. We added the 20 wt% Ca(OH)2 slurry solution (100 ml) into the CO2 saturated solution. This created the Ca2þ ion in the conversion solution which produced the precipitate. Ca(OH)2 occurred in a slurry state in water because of its low solubility. We purchased Ca(OH)2 from Daejung Chemicals & Metals Co., Ltd. Its purity was greater than 95%. After adding the Ca(OH)2, the solution was stirred by a magnetic reactor for 24 h. The formed precipitate was separated from solution using a vacuum pump. We used a glass filter funnel with a porosity of 10e16 mm. The separated solid was dried in an oven at 90 C for 2 days. The dried solid was analyzed using XRD (X-ray diffraction) to check the composition. We measured the residual CO2 in a separate solution using the desorption process. The temperature during this process was maintained by a thermo water bath at 70 C. In this process, in addition to the CO2, only N2 gas was used. Throughout this experiment, the CO2 conversion amounts were calculated. After desorption, we repeated the same experiment in order to check for the possibility of a conversion solution and to determine whether or not the solution could be reused. 3. Results and discussion 3.1. CO2 loading results In this study, we used an MEA solution to convert CO2 to ionic CO2. According to Hook, the reaction of CO2 in an amine solution is simply indicated as follows [21]:
CO2 þ R2 NH4R2 COO NHþ
(1)
R2 COO NHþ þ R2NH4R2 NCOO þ R2 NHþ 2
(2)
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Fig. 1. Schematic diagram of CO2 absorption and conversion experiments.
Considering the MEA and CO2 reaction, we realized that absorbed CO2 changed to carbamate, bicarbonate, and carbonate [22]. Following this mechanism, the amine solution captured the CO2. Similarly, the general CCS process also follows this mechanism, and ionic CO2 was reemitted into an external surface by heat or pressure [23]. After the CO2 was reemitted, the used amine could reabsorb the CO2 via the free amine form [24]. This was one of the main reasons that we selected MEA for our study. The general CO2 loading curve of the absorption-desorption results, according to Seoungmoon Lee et al., is shown in Fig. 2 [3]. The absorbed CO2 in the amine solution curve increased from 0 to approximately 0.54
Fig. 2. General CO2 loading curve of absorption and desorption [3].
when the CO2 was flowed into the amine solution. In the desorption case, a small amount of CO2 gas was reemitted at the start of the desorption experiment [25]. Additionally, absorbed CO2 in the amine solution was rapidly reemitted after a short period of time [26]. However, our study showed a different result compared to previous works. Our results from the 5 and 30 wt% MEA solutions are shown in Figs. 3 and 4, respectively, as indicated by the 1st and 2nd MEA conversion results. The 5 wt% MEA solution was saturated with CO2 for 60 min because its concentration was relatively low. However, the 2nd step experiment of the CO2 absorption has lower CO2 loading values than the results from the 1st experiment. The reasons for this can be explained by the absorption time. As shown in the figure, the total absorption time was different between the 1st and 2nd experiments. Therefore, it appears that the total CO2 loading value decreased when the same experiment was repeated. We believe that the decreased CO2 loading was caused by residual Ca2þ ions in the reused solution. As was done previously, we used a Ca(OH)2 slurry solution. Generally, Ca(OH)2 has a low solubility in water at 25 C. As a result, some Ca(OH)2 remained in the separated solution after the precipitate reaction. In other words, the remaining Ca2þ ions could react with the CO2 flow in the 2nd experiment because some of the particles were generated during the 2nd experiment. Consequentially, the remaining Ca2þ ions could be causing a decrease in the total CO2 loading. This explanation is included in the next section to corroborate our XRD results. The 30 wt% MEA solution showed a different result because it has a similar absorption time [Fig. 4]. This indicates that there is sufficient CO2 to absorb and convert the Ca(OH)2 solution. Therefore, the added Ca(OH)2 solution could react with CO2 dissolved in the MEA solution. We believe that most of the absorbed CO2 in the MEA solution was converted to CaCO3. In addition, we assume that the separated solution consisted primarily of the free amine composition. As a result, the reabsorption time was similar to that
S. Park et al. / Energy 75 (2014) 624e629
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Fig. 3. CO2 concentration of overall process using 5 wt% MEA solution.
of the 1st step. Sangwon Park et al. also obtained similar results [8]. Based on the assumption that sufficient metal ions were provided to the solution, they reported that most of the absorbed CO2 in the amine solution was converted to carbonate [6,8]. Considering Figs. 3 and 4, we can see that residual CO2 did not remain in the MEA solution during the desorption test. This suggests that most of the CO2 was absorbed and converted by the solid particle via the first test. Thus, we checked for residual CO2 in the MEA solution after the desorption test. Moreover, we confirmed the total CO2 loading capacity. This was not significantly different from the 2nd test. We also assumed that desorption in the CCS process could be replaced with CO2 precipitation, suggesting that the absorbents could be reused in the CCS process when used in our process. The separated solution did not reemit the CO2 after the 2nd absorption and conversion. We believe that these results support our assumptions. Similarly, 10 and 30 wt% amine solutions agreed with our results. According to Sangwon Park et al., a 10 wt% amine solution had a lower CO2 loading capacity than a 30 wt% solution [8]. The 10 wt% solution produced a lower amount of carbonate compared to the 30 wt% solution due to the amount of CO2 that was converted. However, our results could differ from their study because they used a 20 wt% CaCl2 solution. CaCl2 solubility in water is greater than the solubility of Ca(OH)2 in water. Accordingly, when
Fig. 4. CO2 concentration of overall process using 30 wt% MEA solution.
Fig. 5. XRD results of formed precipitate by using 5 wt% MEA solution. a) 5 wt% MEA solution and 20 wt% Ca(OH)2 slurry in the first experiment. b) 5 wt% MEA solution and 20 wt% Ca(OH)2 slurry in the second experiment. c: calcite (CaCO3), v: vaterite (CaCO3), a: aragonite (CaCO3), △: calcium hydroxide (Ca(OH)2).
we made a metal additive solution for fixing the ionic CO2, our Ca(OH)2 produced a lower concentration of Ca2þ ions compared to the CaCl2 that was used by Park et al. Consequently, the carbonate formation rate in our study was not significantly faster compared to when the CaCl2 solution was used. 3.2. Verification of precipitate formation by XRD and SEM We separated the solid and liquid by a glass funnel filter after reacting CO2-saturated MEA with the Ca(OH)2 slurry solution. The separated solid was dried in an oven at 90 C for 2 days and analyzed by XRD to check its composition. Results showed that the obtained solid consisted of a large amount of calcium carbonate and some calcium hydroxide. We hypothesized that the calcium carbonate was caused by Ca2þ ions in the Ca(OH)2 slurry and ionic CO2 in the MEA solution. We explain this reaction as follows:
Fig. 6. XRD results of formed precipitate by using 30 wt% MEA solution. a) 30 wt% MEA solution and 20 wt% Ca(OH)2 slurry in the first experiment. b) 30 wt% MEA solution and 20 wt% Ca(OH)2 slurry in the second experiment. c: calcite (CaCO3), v: vaterite (CaCO3), a: aragonite (CaCO3), △: calcium hydroxide (Ca(OH)2).
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Fig. 7. SEM image of formed precipitate. a) 5 wt% MEA solution and 20 wt% Ca(OH)2 slurry in the first experiment. b) 30 wt% MEA solution and 20 wt% Ca(OH)2 slurry in the first experiment.
Ca2þ þ CO2 3 4CaCO3
(3)
2þ In general, 1 mol of CO2 3 reacts with 1 mol of Ca . However, the ionic CO2 does not occur in only a carbonate form [27,28]. According to Hook, CO2 absorbed in amine can be altered by changes in the pH [8]. Therefore, the additional mechanisms described in Equations (4)e(6) are considered.
2RNH2 þ CO2 þ H2 O4RNHCO2 RNH3 þ Ca þ
2þ
/CaCO3 Y
RNHþ 3
(4)
2þ þ RNHþ 3 HCO3 þ RNH2 þ Ca /CaCO3 Y þ 2RNH3
(5)
2 2þ þ 2RNHþ 3 þ CO3 þ Ca /CaCO3 Y þ 2RNH3
(6)
Following Equations (4)e(6), we confirmed that the formed precipitate was carbonate. As shown in Fig. 5, we determined the composition of the precipitate. The formed solid consisted of Ca(OH)2 and CaCO3. First, the formed solid was determined to be Ca(OH)2 based on the characteristic peaks at 18.27, 28.50, 46.73, 51.03, and 54.59 2q, respectively. Calcium carbonate in Fig. 5 is designated as calcite (c), vaterite (v), and aragonite (a). 5 wt% MEA did not sufficiently react with the added Ca2þ when the Ca(OH)2 was measured in a formed solid. The reason for this has already been explained. The amount of absorbed CO2 in 5 wt% MEA remained in solution after forming the solid because the concentration of Ca2þ ions in the Ca(OH)2 slurry was relatively high. Fig. 6 shows that the 30 wt% results were almost nonexistent for Ca(OH)2. As was previously assumed, 30 wt% MEA was sufficient to make calcium carbonate with when Ca2þ ions were added. In addition, a single peak of Ca(OH)2 was measured at 51.03 2q (Fig. 6). This indicates that the Ca(OH)2 has a low solubility in water. This peak was present in both the 1st and 2nd experiments. Therefore, we believe that most of the CO2 in 30 wt% MEA was converted to CaCO3; some Ca(OH)2 was also obtained. In a similar example, Mg(OH)2 reported comparable results [30]. Some of the Mg(OH)2 did not react with CO2 because its concentration was low [31]. We thought that the residual Mg(OH)2 might react with CO2 as previously explained. In conclusion, a high amine concentration could yield a higher amount of carbonate compared to a lower concentration. Although only a small amount of Ca(OH)2 was generated, we believe that Ca(OH)2 had a sufficient source of metal [32]. In addition, we also added to the SEM images in Fig. 7. Fig. 7 (a) and (b) shows that most of the CO2 precipitates that formed were smaller than 2 mm. Also, three different structures were observed: calcite, vaterite, and aragonite. The Ca(OH)2 slurry makes precipitates by reacting with ionic CO2. Nevertheless, the purity is low
because three different structures were produced. We believe that further development is required in the future to increase the purity. However, we studied the basic formation of CO2 fixation by using waste materials as the Ca2þ source. As a result, we checked the possibility of CO2 fixation using Ca(OH)2. 4. Conclusion This study verified the possibility of CO2 fixation by using 5 and 30 wt% MEA solutions with a Ca(OH)2 slurry. The MEA solution rapidly converted CO2 to ionic CO2. Moreover, the Ca(OH)2 slurry converted the Ca2þ ions to ionic CO2. As a result, we can make a precipitate using no additional energy to generate the calcium carbonate. Although 5 wt% MEA has a lower CO2 loading capacity than 30 wt% MEA, we believe that the amine has a sufficient potential to fix CO2 at the point of CO2 transfer velocity [29]. The Ca(OH)2 slurry also yielded enough Ca2þ ions to make carbonate. A stable source of metal ions can fix the converted CO2 in the MEA solution. In addition, the MEA solution can also absorb the CO2 again under the same conditions. Consequently, our previous assumption was confirmed. This study demonstrates how to reduce the additional energy in the CCS process and how to overcome environmental problems associated with construction waste including alkalinity waste, leachate, etc. We are planning to develop the Ca(OH)2 solubility in order to obtain many more Ca2þ ions by controlling the physical conditions or by using additives. Also, we plan to use a longer experiment times to verify the efficacy of long-term absorbent reuse. We believe that this study is not only applicable for reusing emitted CO2 but also for reusing construction waste. Acknowledgments This work was supported by a National Research Foundation of Korea (NRF) grant funded by the Korean government (MEST) (No. 2011-0029161). This work was also supported by a Korea Agency for Infrastructure Technology Advancement (KAIA) grant funded by the Korean government (MOLIT) (No. 13 construction reasearch R02). References [1] IPCC. In: Metz B, Davidson O, de Coninck HC, Loos M, Meyer LA, editors. IPCC special report on carbon dioxide capture and storage. Prepared by working group III of the intergovernmental panel on climate change. Cambridge (United Kingdom)/New York (NY, USA): Cambridge University Press; 2005. p. 442. [2] IEA. Energy technology perspectives: BLUE map scenario. OECD/IEA; 2011.
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Energy journal homepage: www.elsevier.com/locate/energy
A study of CO2 precipitation method considering an ionic CO2 and Ca(OH)2 slurry Sangwon Park, Hoyong Jo, Dongwoo Kang, Jinwon Park* Department of Chemical & Biomolecular Engineering, Yonsei University, 50 Yonsei-ro, Seodaemun-gu, Seoul 120-749, Republic of Korea
a r t i c l e i n f o
a b s t r a c t
Article history: Received 18 January 2014 Received in revised form 2 July 2014 Accepted 9 August 2014 Available online 28 August 2014
CCS (carbon capture and storage) is the most popular technology used for the reduction of CO2 in the post-combustion stage. However, the CCS process has some disadvantages including uncertainty about the stability of the land that is used to store the separated CO2. Consequently, CCU (carbon capture and utilization) technologies have recently received increased attention as a possible replacement for CCS. In this study, we utilized CO2 fixation methods by using the metal carbonate mechanism. We selected 5 and 30 wt% MEA (mono-ethanolamine) solutions to rapidly make a carbonate and Ca(OH)2 slurry. In all of the experiments, normal temperature and pressure conditions were maintained (except during desorption to check for residual CO2 in the MEA solution). Consequently, most of the CO2 was converted to carbonate. The MEA converted CO2 to ionic CO2 and rapidly created calcium carbonate. Also the formed solids that were observed were determined to be CaCO3 and Ca(OH)2 by X-ray diffractometry. Also, the MEA solution could be reused to absorb CO2. Therefore, we have confirmed the development of our suggested CCS process. This process has the ability not only to reuse emitted CO2, but it can also be employed to reuse construction wastes that include heavy metals. © 2014 Elsevier Ltd. All rights reserved.
Keywords: CCS (carbon capture and storage) process CO2 mineralization PCC (precipitate calcium carbonate)
1. Introduction It is widely accepted that GHGs (greenhouse gases) are responsible for the greenhouse effect [1]. Most GHGs are emitted by human activities such as from coal-fired power plants that use fossil fuels [2]. Thus, many researchers have been exploring various CO2 reduction methods to curtail the greenhouse effect. The CCS (carbon capture and storage) process is a well-known technology that is widely employed to reduce CO2 emissions after combustion [3,4]. CCS appears particularly well-suited for the relatively large, high concentration CO2 emissions found in many types of industry [4]. For example, CCS technology research has focused on coal-fired power plants in Korea. However, the CCS process has some disadvantages. First, a large area is required to separate CO2 from the absorbents. As previously stated, it is beneficial to apply CCS technologies to industries that produce high CO2 emissions because the amount of CO2 that is separated is equal to the amount that is absorbed using this process. According to the KEPRI (Korea Electric Power Research Institute), a 10 MW coal-fired power plant emits 80,000 tons of CO2 in a single year [5]. A large area of land would be
* Corresponding author. Tel.: þ82 2 364 1807; fax: þ82 2 312 6401. E-mail address: [email protected] (J. Park). http://dx.doi.org/10.1016/j.energy.2014.08.036 0360-5442/© 2014 Elsevier Ltd. All rights reserved.
required to store this CO2. This poses a problem for countries such as Korea, which have a small total land area. Secondly, there are issues regarding the stability of storing CO2 in the land or ocean [6,7]. Separated CO2 is stored underground or in the ocean using absorbents. However, there is the possibility that the stored CO2 could be reemitted into the atmosphere after physical changes to the storage area such as damage caused from earthquakes or volcanoes. Finally, there are also economic problems. The CCS process is made up of two parts: absorption and desorption [8]. Between these, CO2 desorption uses the majority of the energy (more than 70%) [6]. In the desorption process, CO2 in the absorbents is separated using heat or pressure. As a result, the CCS process consumes a large amount of energy [9]. Therefore, some researchers have attempted to find ways to reduce energy consumption during desorption [10,11]. We believe that if certain improvements are made (e.g. researching alternative absorbents and additives to reduce energy consumption), CCS can become superior to other CO2 reduction methods from an economic standpoint [12]. CCU (carbon capture and utilization) technologies have been studied since the 1990s as a potential method to overcome the aforementioned shortcomings of CCS [13]. The goal of CCU is to reuse captured CO2 via biological, physical, and/or chemical means. In the biological case, captured CO2 can be converted into renewable energy or acetic acid using algae [6,10]. In the physical and
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chemical cases, CO2 fixation and reuse methods were used. These methods focused on the chemical conversion mechanism. For example, CO2 gas was converted to a solid (e.g. calcium, barium, or magnesium carbonate) by combining it with metal ions [14,15]. Solid CO2 conversion byproducts have better chemical stability compared to gases or ionic liquids. Following these methods, the emitted CO2 can be stored underground or in the ocean and ultimately reused [6]. With regard to the metal carbonate mechanism, CO2 fixation methods were classified as using either high- or lowenergy depending on the total energy consumption. In general, high-energy consumption methods were used to directly emit CO2 and some of the metals [16,17]. These methods require high temperature and pressure to make the metal carbonate because metal ions, including metal materials, need a certain amount of activation energy to react with CO2 gas. According to Lackner et al., calcium and magnesium carbonates from serpentine require temperatures above 500 C and a pressure greater than 340 bars [18]. In addition, magnesium carbonate from magnesium hydroxide requires 540 C and 45 bars [19]. These methods appear to have a significant disadvantage from the standpoint of energy consumption. However, some industries, including power plants and the cement industry, do have sufficient energy to support high-temperature fixation methods. In addition, the waste heat generated in this process could be used [10]. Conversely, low-energy CO2 fixation does not require high temperature or high pressure [6,10]. The methods mentioned above use conversion materials. In other words, emitted CO2 is converted to ionic CO2 by conversion materials such as amine, ammonia, and alkaline solutions, which are well-known CO2 absorbents [9e11]. Acid gas CO2 is converted to ionic CO2 during absorption in an alkaline solution. An alkaline solution generally creates ionic CO2 via CO2 absorption, carbonate, bicarbonate, and carbamate, respectively. Then, the ionic CO2 reacts with metal ions. Therefore, ionic CO2 could form a metal carbonate more easily than gaseous CO2 because it has a lower activation energy. According to Sangwon Park et al., low-energy consumption CO2 fixation methods are superior to the CCS process because they maintain normal conditions of 30 C and 1 bar. They also do not require a desorption process [6,10]. In addition, most of the CO2 absorbed in their amine solution was converted to metal carbonates. This means that additional energy was not required in desorption and regeneration of the amine solution. Therefore, we believe that this method is preferable from an economic standpoint. However, when considering the aforementioned studies, we believe that two important points should be emphasized. First, the rate at which CO2 is converted to ionic CO2 and the absorbed CO2 conversion kinetics must be taken into account. Second, it must be determined whether or not the fixation sources can provide a sufficient amount of ionic CO2 to produce a metal carbonate. A limited supply of metal ions might not produce enough metal carbonate. In other words, converted CO2 would remain in solution because the total CO2 loading is larger than the amount of supplied metal ions. In this regard, this method is similar to CCS. Additionally, a low CO2 conversion rate could incur a larger reactor size when enough metal is provided. This means that additional energy consumption would be required (similar to the high-energy consumption process for rapidly making metal carbonates in smaller reactors). Thus, we have taken these facts into consideration for CO2 fixation under low-energy consumption conditions. We explored ways to make a carbonate using CCU under normal temperatures and pressures. We selected an amine solution in order to convert CO2, and we selected Ca(OH)2 as the metal source. We chose the amine solution because it has a high CO2 conversion rate. In addition, we assumed that this study could improve CCS by considering desorption; thus, we determined that an amine solution would make a suitable conversion material. Finally, we
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selected MEA (mono-ethanolamine) because it was widely used in most CCS processes and has been verified by several researchers [3e5,20]. Ca(OH)2 was easily obtained from concrete and construction wastes. Therefore, this study verified the possibility of simultaneously obtaining CO2 and Ca(OH)2 emitted from construction waste under normal conditions. 2. Experimental The experimental apparatus is shown in Fig. 1. We removed the remaining CO2 by using high-purity N2 gas (99.999%). The highpurity N2 gas caught the residual CO2 in the experiment, and the mixing gas flowed through a Teflon tube. The mixed gases were detected with a CO2 analyzer at 30 s intervals. We maintained this process for 30 min. After that, we assumed that the CO2 concentration was 15 V/V% in a balance of N2. N2 and CO2 gases were controlled by MFC (mass flow rate), and the flow rates were 1870 and 300 ml/min, respectively. Also, the purity of the CO2 and N2 gases used in the absorption-desorption process were both 99.9%. Each gas was mixed by a saturated reactor which flowed into the conversion reactor. The conversion reactor was filled with 5 and 30 wt% MEA solutions. Each MEA solution was 400 ml and balanced on water using the corresponding weight ratio. MEA was purchased from SigmaeAldrich and was more than 99% pure. Simulated gas continuously flowed into the CO2 analyzer, and its concentration was analyzed at 30 s intervals. After the conversion reaction, the simulated gas was flowed through a cold condenser (15 C). This was used to maintain the concentration of the conversion solution via trapping any water loss. A reaction temperature of 30 C was maintained in the saturated and conversion reactors by a thermo water bath. Finally, we assumed that the conversion experiment was finished when the outlet CO2 concentration reached 15 V/V%. After the conversion experiment was completed, the CO2 precipitation experiment began. We added the 20 wt% Ca(OH)2 slurry solution (100 ml) into the CO2 saturated solution. This created the Ca2þ ion in the conversion solution which produced the precipitate. Ca(OH)2 occurred in a slurry state in water because of its low solubility. We purchased Ca(OH)2 from Daejung Chemicals & Metals Co., Ltd. Its purity was greater than 95%. After adding the Ca(OH)2, the solution was stirred by a magnetic reactor for 24 h. The formed precipitate was separated from solution using a vacuum pump. We used a glass filter funnel with a porosity of 10e16 mm. The separated solid was dried in an oven at 90 C for 2 days. The dried solid was analyzed using XRD (X-ray diffraction) to check the composition. We measured the residual CO2 in a separate solution using the desorption process. The temperature during this process was maintained by a thermo water bath at 70 C. In this process, in addition to the CO2, only N2 gas was used. Throughout this experiment, the CO2 conversion amounts were calculated. After desorption, we repeated the same experiment in order to check for the possibility of a conversion solution and to determine whether or not the solution could be reused. 3. Results and discussion 3.1. CO2 loading results In this study, we used an MEA solution to convert CO2 to ionic CO2. According to Hook, the reaction of CO2 in an amine solution is simply indicated as follows [21]:
CO2 þ R2 NH4R2 COO NHþ
(1)
R2 COO NHþ þ R2NH4R2 NCOO þ R2 NHþ 2
(2)
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Fig. 1. Schematic diagram of CO2 absorption and conversion experiments.
Considering the MEA and CO2 reaction, we realized that absorbed CO2 changed to carbamate, bicarbonate, and carbonate [22]. Following this mechanism, the amine solution captured the CO2. Similarly, the general CCS process also follows this mechanism, and ionic CO2 was reemitted into an external surface by heat or pressure [23]. After the CO2 was reemitted, the used amine could reabsorb the CO2 via the free amine form [24]. This was one of the main reasons that we selected MEA for our study. The general CO2 loading curve of the absorption-desorption results, according to Seoungmoon Lee et al., is shown in Fig. 2 [3]. The absorbed CO2 in the amine solution curve increased from 0 to approximately 0.54
Fig. 2. General CO2 loading curve of absorption and desorption [3].
when the CO2 was flowed into the amine solution. In the desorption case, a small amount of CO2 gas was reemitted at the start of the desorption experiment [25]. Additionally, absorbed CO2 in the amine solution was rapidly reemitted after a short period of time [26]. However, our study showed a different result compared to previous works. Our results from the 5 and 30 wt% MEA solutions are shown in Figs. 3 and 4, respectively, as indicated by the 1st and 2nd MEA conversion results. The 5 wt% MEA solution was saturated with CO2 for 60 min because its concentration was relatively low. However, the 2nd step experiment of the CO2 absorption has lower CO2 loading values than the results from the 1st experiment. The reasons for this can be explained by the absorption time. As shown in the figure, the total absorption time was different between the 1st and 2nd experiments. Therefore, it appears that the total CO2 loading value decreased when the same experiment was repeated. We believe that the decreased CO2 loading was caused by residual Ca2þ ions in the reused solution. As was done previously, we used a Ca(OH)2 slurry solution. Generally, Ca(OH)2 has a low solubility in water at 25 C. As a result, some Ca(OH)2 remained in the separated solution after the precipitate reaction. In other words, the remaining Ca2þ ions could react with the CO2 flow in the 2nd experiment because some of the particles were generated during the 2nd experiment. Consequentially, the remaining Ca2þ ions could be causing a decrease in the total CO2 loading. This explanation is included in the next section to corroborate our XRD results. The 30 wt% MEA solution showed a different result because it has a similar absorption time [Fig. 4]. This indicates that there is sufficient CO2 to absorb and convert the Ca(OH)2 solution. Therefore, the added Ca(OH)2 solution could react with CO2 dissolved in the MEA solution. We believe that most of the absorbed CO2 in the MEA solution was converted to CaCO3. In addition, we assume that the separated solution consisted primarily of the free amine composition. As a result, the reabsorption time was similar to that
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Fig. 3. CO2 concentration of overall process using 5 wt% MEA solution.
of the 1st step. Sangwon Park et al. also obtained similar results [8]. Based on the assumption that sufficient metal ions were provided to the solution, they reported that most of the absorbed CO2 in the amine solution was converted to carbonate [6,8]. Considering Figs. 3 and 4, we can see that residual CO2 did not remain in the MEA solution during the desorption test. This suggests that most of the CO2 was absorbed and converted by the solid particle via the first test. Thus, we checked for residual CO2 in the MEA solution after the desorption test. Moreover, we confirmed the total CO2 loading capacity. This was not significantly different from the 2nd test. We also assumed that desorption in the CCS process could be replaced with CO2 precipitation, suggesting that the absorbents could be reused in the CCS process when used in our process. The separated solution did not reemit the CO2 after the 2nd absorption and conversion. We believe that these results support our assumptions. Similarly, 10 and 30 wt% amine solutions agreed with our results. According to Sangwon Park et al., a 10 wt% amine solution had a lower CO2 loading capacity than a 30 wt% solution [8]. The 10 wt% solution produced a lower amount of carbonate compared to the 30 wt% solution due to the amount of CO2 that was converted. However, our results could differ from their study because they used a 20 wt% CaCl2 solution. CaCl2 solubility in water is greater than the solubility of Ca(OH)2 in water. Accordingly, when
Fig. 4. CO2 concentration of overall process using 30 wt% MEA solution.
Fig. 5. XRD results of formed precipitate by using 5 wt% MEA solution. a) 5 wt% MEA solution and 20 wt% Ca(OH)2 slurry in the first experiment. b) 5 wt% MEA solution and 20 wt% Ca(OH)2 slurry in the second experiment. c: calcite (CaCO3), v: vaterite (CaCO3), a: aragonite (CaCO3), △: calcium hydroxide (Ca(OH)2).
we made a metal additive solution for fixing the ionic CO2, our Ca(OH)2 produced a lower concentration of Ca2þ ions compared to the CaCl2 that was used by Park et al. Consequently, the carbonate formation rate in our study was not significantly faster compared to when the CaCl2 solution was used. 3.2. Verification of precipitate formation by XRD and SEM We separated the solid and liquid by a glass funnel filter after reacting CO2-saturated MEA with the Ca(OH)2 slurry solution. The separated solid was dried in an oven at 90 C for 2 days and analyzed by XRD to check its composition. Results showed that the obtained solid consisted of a large amount of calcium carbonate and some calcium hydroxide. We hypothesized that the calcium carbonate was caused by Ca2þ ions in the Ca(OH)2 slurry and ionic CO2 in the MEA solution. We explain this reaction as follows:
Fig. 6. XRD results of formed precipitate by using 30 wt% MEA solution. a) 30 wt% MEA solution and 20 wt% Ca(OH)2 slurry in the first experiment. b) 30 wt% MEA solution and 20 wt% Ca(OH)2 slurry in the second experiment. c: calcite (CaCO3), v: vaterite (CaCO3), a: aragonite (CaCO3), △: calcium hydroxide (Ca(OH)2).
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Fig. 7. SEM image of formed precipitate. a) 5 wt% MEA solution and 20 wt% Ca(OH)2 slurry in the first experiment. b) 30 wt% MEA solution and 20 wt% Ca(OH)2 slurry in the first experiment.
Ca2þ þ CO2 3 4CaCO3
(3)
2þ In general, 1 mol of CO2 3 reacts with 1 mol of Ca . However, the ionic CO2 does not occur in only a carbonate form [27,28]. According to Hook, CO2 absorbed in amine can be altered by changes in the pH [8]. Therefore, the additional mechanisms described in Equations (4)e(6) are considered.
2RNH2 þ CO2 þ H2 O4RNHCO2 RNH3 þ Ca þ
2þ
/CaCO3 Y
RNHþ 3
(4)
2þ þ RNHþ 3 HCO3 þ RNH2 þ Ca /CaCO3 Y þ 2RNH3
(5)
2 2þ þ 2RNHþ 3 þ CO3 þ Ca /CaCO3 Y þ 2RNH3
(6)
Following Equations (4)e(6), we confirmed that the formed precipitate was carbonate. As shown in Fig. 5, we determined the composition of the precipitate. The formed solid consisted of Ca(OH)2 and CaCO3. First, the formed solid was determined to be Ca(OH)2 based on the characteristic peaks at 18.27, 28.50, 46.73, 51.03, and 54.59 2q, respectively. Calcium carbonate in Fig. 5 is designated as calcite (c), vaterite (v), and aragonite (a). 5 wt% MEA did not sufficiently react with the added Ca2þ when the Ca(OH)2 was measured in a formed solid. The reason for this has already been explained. The amount of absorbed CO2 in 5 wt% MEA remained in solution after forming the solid because the concentration of Ca2þ ions in the Ca(OH)2 slurry was relatively high. Fig. 6 shows that the 30 wt% results were almost nonexistent for Ca(OH)2. As was previously assumed, 30 wt% MEA was sufficient to make calcium carbonate with when Ca2þ ions were added. In addition, a single peak of Ca(OH)2 was measured at 51.03 2q (Fig. 6). This indicates that the Ca(OH)2 has a low solubility in water. This peak was present in both the 1st and 2nd experiments. Therefore, we believe that most of the CO2 in 30 wt% MEA was converted to CaCO3; some Ca(OH)2 was also obtained. In a similar example, Mg(OH)2 reported comparable results [30]. Some of the Mg(OH)2 did not react with CO2 because its concentration was low [31]. We thought that the residual Mg(OH)2 might react with CO2 as previously explained. In conclusion, a high amine concentration could yield a higher amount of carbonate compared to a lower concentration. Although only a small amount of Ca(OH)2 was generated, we believe that Ca(OH)2 had a sufficient source of metal [32]. In addition, we also added to the SEM images in Fig. 7. Fig. 7 (a) and (b) shows that most of the CO2 precipitates that formed were smaller than 2 mm. Also, three different structures were observed: calcite, vaterite, and aragonite. The Ca(OH)2 slurry makes precipitates by reacting with ionic CO2. Nevertheless, the purity is low
because three different structures were produced. We believe that further development is required in the future to increase the purity. However, we studied the basic formation of CO2 fixation by using waste materials as the Ca2þ source. As a result, we checked the possibility of CO2 fixation using Ca(OH)2. 4. Conclusion This study verified the possibility of CO2 fixation by using 5 and 30 wt% MEA solutions with a Ca(OH)2 slurry. The MEA solution rapidly converted CO2 to ionic CO2. Moreover, the Ca(OH)2 slurry converted the Ca2þ ions to ionic CO2. As a result, we can make a precipitate using no additional energy to generate the calcium carbonate. Although 5 wt% MEA has a lower CO2 loading capacity than 30 wt% MEA, we believe that the amine has a sufficient potential to fix CO2 at the point of CO2 transfer velocity [29]. The Ca(OH)2 slurry also yielded enough Ca2þ ions to make carbonate. A stable source of metal ions can fix the converted CO2 in the MEA solution. In addition, the MEA solution can also absorb the CO2 again under the same conditions. Consequently, our previous assumption was confirmed. This study demonstrates how to reduce the additional energy in the CCS process and how to overcome environmental problems associated with construction waste including alkalinity waste, leachate, etc. We are planning to develop the Ca(OH)2 solubility in order to obtain many more Ca2þ ions by controlling the physical conditions or by using additives. Also, we plan to use a longer experiment times to verify the efficacy of long-term absorbent reuse. We believe that this study is not only applicable for reusing emitted CO2 but also for reusing construction waste. Acknowledgments This work was supported by a National Research Foundation of Korea (NRF) grant funded by the Korean government (MEST) (No. 2011-0029161). This work was also supported by a Korea Agency for Infrastructure Technology Advancement (KAIA) grant funded by the Korean government (MOLIT) (No. 13 construction reasearch R02). References [1] IPCC. In: Metz B, Davidson O, de Coninck HC, Loos M, Meyer LA, editors. IPCC special report on carbon dioxide capture and storage. Prepared by working group III of the intergovernmental panel on climate change. Cambridge (United Kingdom)/New York (NY, USA): Cambridge University Press; 2005. p. 442. [2] IEA. Energy technology perspectives: BLUE map scenario. OECD/IEA; 2011.
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