ARTICLE pubs.acs.org/IECR
CO2 Fixation in Ca2þ-/Mg2þ-Rich Aqueous Solutions through Enhanced Carbonate Precipitation Wenlong Wang,* Mingqiang Hu, Yanli Zheng, Peng Wang, and Chunyuan Ma National Engineering Laboratory for Coal-Fired Pollutants Emission Reduction, Shandong University, Jinan, China 250061 ABSTRACT: In this work, the possibility of achieving fixation of CO2 using Ca and Mg ions was tested and verified. Concentrated seawater from desalination plants, subsurface brines, industrial effluents with high hardness, and/or natural seawaters that are rich in Ca2þ and Mg2þ could all be potential aqueous sources. Theoretical analyses indicated that the carbonation reaction could be enhanced by raising the pH or the CO2 partial pressure. Experiments using synthesized seawater confirmed this possibility. Over 90% of the Ca2þ and Mg2þ ions in the seawater could be converted by precipitation in the forms of MgCO3 and dolomite [MgCa(CO3)2], and the kinetics of the process was found to be quite acceptable. It was found that 1 m3 of natural seawater could fix about 1.34 m3 or 2.65 kg of CO2 (gas volume, standard conditions), and the potential of concentrated seawater is 23 times this value. Even if the annual CO2 emissions of the entire world were captured in this way, the concentration of Ca2þ/Mg2þ in natural seawater would change at only the part-per-million scale, such that the ecological effects could be negligible. This idea has great potential for application. It might be able to realize not only the permanent fixation of CO2 but also the production of large amounts of carbonate byproducts.
1. INTRODUCTION With increasing concern about human-induced global climate change, CO2 becomes a primary focus because its increasing concentration is believed to be the most important factor in the greenhouse effect.1 Resorting to carbon-free or carbon-neutral energy sources and developing a so-called low-carbon economy are unanimously admitted to be radical solutions. Carbon capture and sequestration (CCS) is another. Many technologies and ideas have been developed for CCS: Typical separation methods include chemical or physical absorption/adsorption, membrane systems, cryogenic fractionation, chemical looping combustion, and so on,27 and sequestration is focused on storage in deep geological formations, in deep ocean masses, or in the form of mineral carbonates.817 However, it is certain that no single technology can solve all of the problems. Novel methods will continue to be proposed in the future. Mineralization through reactions between CO2 and magnesium/calcium silicate rocks can be regarded as a type of permanent fixation for CO2. This work suggests another approach to fixation. Because concentrated seawater from desalination plants, subsurface brines, industrial effluents with high hardness, and/or natural seawaters are all rich in Ca2þ and Mg2þ ions, our idea is to directly enhance carbonate precipitation to achieve CO2 fixation with these aqueous sources. Because these sources are different only in their concentrations of ions, we consider natural seawater, the most common source, as the research object. In fact, researchers are showing increasing interests in carbonate formation and dissolution in the ocean because of the role of these reactions in the CO2 issue. Morse et al. and Millero have published careful reviews on this topic.18,19 However, most of the previous work has focused mainly on carbonation reactions under natural conditions and from a biological perspective.2022 The possibility of CO2 fixation by intentionally enhancing the carbonate precipitation process is still unknown. Therefore, in this work, we investigated the possibility of this idea through calculations and experiments to obtain a better understanding of Mg2þ/Ca2þ utilization. r 2011 American Chemical Society
Seawater contains more than 80 types of chemical elements, most of which appear as ions. Among them, the contents of 10 ions can reach the level of 1 ppm. They are called “main components” and make up more than 99% of the total salts in seawater. Table 1 shows the detailed ionic contents of a typical sample of seawater. It can be seen that even the top 10 ions vary significantly in content. Among cations, Mg2þ and Ca2þ rank second and third, respectively, and the dissolved inorganic carbon (H2CO3*, HCO3, and CO32) mainly exists in the form of HCO3. Normally, the formation of precipitation products such as calcium or magnesium carbonate is difficult in seawater, because these carbonates are far below saturation limits in the water system mostly because of the low concentration of CO32. Therefore, our aim is to change the ionic concentrations and saturation degrees to enhance the precipitation process.
2. FEASIBILITY 2.1. Solubility. As is well-known, all compounds are soluble in water to a certain extent. The most general form of a precipitationdissolution reaction is
A a Bb T aA xþ þ bBy
ð1Þ
When the background electrolyte concentration is high (e.g., in seawater, landfill leachates, and activated sludge), the solubility product constant can be defined as Ks ¼ fA xþ ga fBy gb ¼ ðγA ½A xþ Þa ðγB ½By Þb
ð2Þ
where {i} represents the activity of species i, γi is its activity coefficient, and [i] is its molar concentration. The Ks values of Received: December 20, 2010 Accepted: May 19, 2011 Revised: April 14, 2011 Published: May 19, 2011 8333
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Table 1. Contents of 10 Main Types of Ions in Seawater23 ion
mass concentration (mg/L)
molar concentration m (mol/L)
ion
468.6
Cl SO42
Naþ
10773
Mg2þ
1294
53.24
Ca2þ
412
Kþ Sr2þ
399 8
10.28 10.21
Br
0.09
F
Ksa
compound
4.95 109
CaCO3
5
3.73 10
CaSO4
Ksa
MgCO3
1.15 105
Mg(OH)2
5.66 1012
6
7.88 10
Ca(OH)2 a
Note: Ks depends on temperature only and the values of K at different temperatures can be determined by thermodynamic calculations.
some slightly or sparingly soluble calcium/magnesium compounds are listed in Table 2. Activity coefficients can be determined by a number of approximations. The Davies equation is presented here because it is valid for the widest range of electrolyte concentrations (where ionic strength, I, is less than 0.5 mol/L).25 The Davies equation is expressed as ! pffiffi I pffiffi 0:2I ð3Þ log γ ¼ AZ2 1þ I
∑
1 C i Zi 2 2
where Ci and Zi are the molar concentration and charge, respectively, of the ith ion in solution. If Qc = [Axþ]a[By]b is employed to describe the actual power product of ion concentrations, then Qc will equal the constant Ks when precipitationdissolution equilibrium is reached in a solution system. For Qc > Ks, the system is supersaturated, and the ionic solid will precipitate from the solution until precipitation dissolution equilibrium. For Qc < Ks, the system is undersaturated, and more solids could be dissolved in it. The precipitationdissolution equilibrium equations for calcium and magnesium carbonates are given by Ca þ CO3 T CaCO3 ðsÞ ¼ fCa2þ gfCO3 2 g
Ks ðCaCO3 Þ
Mg2þ þ CO3 2 T MgCO3 ðsÞ ¼ fMg2þ gfCO3 2 g
Ks ðMgCO3 Þ
2þ
2
ð4Þ
28.23
142
2.33
67
0.84
1.3
0.07
unsaturation. Although precipitation is inhibited even in some cases of supersaturation because of the complicated combination of factors in the ocean,18 this phenomenon is beyond the scope of our discussion. Generally, to accelerate the carbonate precipitation reactions, the Qc values of calcium and magnesium carbonates must be increased. 2.2. Methods of Reaction Enhancement. There are only two ways to increase the values of Qc for CaCO3 and MgCO3: to increase the concentration of Ca2þ/Mg2þ or to increase the concentration of CO32. The concentrations of Ca2þ and Mg2þ in natural seawater are relatively stable. However, their values in concentrated seawater from desalination plants are 23 times higher than those in natural seawater.26 Therefore, combining CO2 fixation with seawater desalination might provide a potential approach. Yet, there are more possible ways of changing the Qc value of CaCO3 or MgCO3 than by increasing the concentration of CO32. In fact, seawater is a natural carbon buffer system, which makes the ocean an enormous sink for CO2. The natural buffer is produced by the following reactions
CO2 ðgÞ þ H2 O T H2 CO3 T Hþ þ HCO3 T 2Hþ þ CO3 2
ð6Þ
The overall process can be divided into several steps. The dissolved CO2 interacts with ambient CO2 with a local equilibrium defined by Henry’s law. The component denoted as H2CO3* includes true carbonic acid and dissolved CO2 CO2 ðgÞ þ H2 O T H2 CO3
½H2 CO3 ¼ KH pCO2 ð7Þ
Carbonic acid dissociates, and the equilibrium concentrations of bicarbonate and carbonate are determined by the first and second dissociation constants H2 CO3 T Hþ þ HCO3
Ka1 ¼
½Hþ ½HCO3 ð8Þ ½H2 CO3
HCO3 T Hþ þ CO3 2
Ka2 ¼
½Hþ ½CO3 2 ½HCO3
ð9Þ
Then, the concentration of CO32 can be determined as ½CO3 2 ¼
ð5Þ
Based on the concentrations of Ca and Mg ions listed in Table 1 and the Ks values listed in Table 2, the saturated concentrations of carbonate ion (CO32) when CaCO3 and MgCO3 begin to precipitate in natural seawater at 298 K can be calculated as 1.55 105 and 8.5 103 mol/L, respectively. In bulk natural seawater, however, the CO32 concentration is only about 2.5 106 mol/L, and calcium and magnesium carbonates are generally in the state of
545.6
2712
where A = 0.5 for water at 25 °C, Z is the charge of the ion whose activity is being calculated, and I is given by I ¼
molar concentration m (mol/L)
19344
HCO3
Table 2. Solubility Product Constants of Some Insoluble Strong Electrolytes of Calcium and Magnesium (298 K)24 compound
mass concentration (mg/L)
Ka1 Ka2 KH pCO2 ½Hþ 2
ð10Þ
or log ½CO3 2 ¼ logðKa1 Ka2 KH pCO2 Þ þ 2pH
ð11Þ
Because Ka1, Ka2, and KH are all constants at a certain temperature, the concentration of CO32 will change only as the atmospheric CO2 partial pressure or pH changes. Therefore, it is feasible to increase the CO32 concentration either by enhancing 8334
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Figure 1. Effects of CO2 partial pressure on CO32 concentration at pH 8: (A) precipitation critical point of CaCO3 (0.00219 atm, 1.55 105 mol/L), (B) precipitation critical point of MgCO3 (1.20 atm, 8.5 103 mol/L).
Figure 2. Effects of pH and CO2 partial pressure on CO32 concentration.
the atmospheric CO2 partial pressure or by increasing the pH of seawater. 2.3. Effects of CO2 Partial Pressure and pH. The current atmospheric CO2 concentration is about 390 ppm. The corresponding partial pressure is 0.00039 atm, a level at which not much CO2 is dissolved in water. However, the concentration of CO2 in waste gases from carbon-intensive plants is usually several orders of magnitude higher than that in the atmosphere. For instance, the partial pressure of CO2 in flue gas exhaust from coalfired power plants is about 0.15 atm. In addition, the partial pressure can be even much higher when CO2 is stored in deep geological formations.17 Therefore, the effects of CO2 pressure on the concentration of CO32 were calculated over a wide pressure range, as shown in Figures 1 and 2. It can be seen in Figure 1 that, at a certain pH, the concentration of CO32 increases linearly with increasing CO2 partial pressure. At pH 8, the CO32 concentration can reach the precipitation formation limit of CaCO3 (i.e., 1.55 105 mol/L) as long as the CO2 partial pressure is 0.00219 atm. Accordingly, it is entirely possible for CaCO3 precipitation to occur at a CO2 partial pressure of 0.15 atm, but this is still not sufficient for MgCO3, whose formation requires a theoretical pCO2 value of 1.2 atm. However, Figure 2 indicates that the CO32 concentration is only near 1 mol/L (i.e., log [CO32] = 0) even when the pCO2 reaches 300 atm at pH 8. This means that the effect of the CO2 partial pressure itself is limited from this point of view. Considering that the CO2 partial pressure is usually not that high, our study focused on the case of 0.15 atm.
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The effects of different pH values are also shown in Figure 2. With the use of a logarithmic scale, the changes in CO32 concentration are illustrated for a wide range of pCO2 values. Obviously, the effect of pH is dramatic; higher pH leads to higher CO32 concentration. When the pH rises from 8 to 9 or from 9 to 10, the CO32 concentration increases by approximately a factor of 102 = 100 at any certain pCO2 value. For instance, at the same pCO2 value of 0.15 atm, the concentration of CO32 is 1.06 103 mol/L at pH 8, whereas it becomes 0.106 mol/L at pH 9. Under such conditions, the precipitation reactions of both CaCO3 and MgCO3 will quite possibly be thermodynamically favorable. Thus, increasing the pH is a reasonable approach to enhancing the precipitation of carbonate. Nevertheless, in practical applications, control of the pH might be one of the most important steps. Ammonia/ammonium chloride buffer can be employed to keep the pH relatively stable in experiments.27 For commercial operations, we are currently carrying out research on increasing pH with ammonia or amine solutions and on regenerating the materials through membrane separation or thermal separation,28 and further results will be reported in the future. In addition, Soong et al. added CaO-rich fly ashes to brine solutions to increase the pH, and House et al. proposed some processes to increase the pH of seawater by electrolytically removing hydrochloric acid from the ocean.29,30 In commercial process designs, these ideas might also be employed. 2.4. Effects of Other Precipitation Products. CaCO3 and MgCO3 are our target products, although other possible precipitation products, such as CaSO4, Ca(OH)2, and Mg(OH)2, should be examined carefully to avoid their formation in the process of increasing CO32 concentration by raising the pH or the partial pressure of CO2. Comparing Table 3 with Table 2, one can see that, in the pH range of 810 and at the CO2 partial pressure of 0.15 atm, the Qc values of CaCO3 and MgCO3 are mostly higher than their Ks values, but those of CaSO4, Ca(OH)2, and Mg(OH)2 are mostly lower than their Ks values, except for that of Mg(OH)2 at pH 10. This means that Mg(OH)2 might precipitate when the pH is too high, in which case unexpected consumption of magnesium and hydroxyl ions would occur. It can be determined that the saturated concentration of hydroxyl at which Mg(OH)2 begins to precipitate is 5.33 105 mol/L, equivalent to pH 9.73. Therefore, this pH limit should not be ignored in any attempts to enhance carbonate precipitation.
3. EXPERIMENTS The above analyses indicate that CO2 fixation in seawater through enhanced carbonate precipitation is theoretically feasible. Further understanding must be based on experiments. Thus, experiments were carried out to explore the practical process of precipitate formation. Figure 3 presents a schematic diagram of the experimental setup. Synthetic gas (15 vol. % CO2 and 85 vol. % N2) was released from a gas cylinder and bubbled through the seawater solution. The seawater reactor was placed in a water bath, which could provide a constant reaction temperature from ambient temperature to 100 °C. A magnetic stirrer was employed throughout the process to obtain a well-distributed reaction mixture. As the reaction solution, artificial seawater was prepared with analytical-grade reagents, whose contents are listed in Table 4. The reaction temperature was set to 25 °C, and the gas flow rate was controlled at 300 mL/min. Because the reactor was an 8335
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Table 3. Qc Values of Possible Precipitation Products under Different Conditions Qc pH
pCO2 (atm)
8 9 10
[OH] (mol/L)
[CO32] (mol/L)
6
0.15
1.06 10
10
0.15
5
1.06 10
4
1.06 10
10
0.15
3
10
CaCO3
MgCO3
7
5.97 10
1
5
5.97 10
3
5.97 10
1
3.08 10
Ca(OH)2
6
3.08 10
4
3.08 10
2
1.16 10
15 13
1.16 10 1.16 10
11
Mg(OH)2
CaSO4
15
1.59 105
13
1.59 105
11
1.59 105
6.02 10 6.02 10
6.02 10
Figure 3. Experimental setup for the carbonation reaction between gaseous CO2 and seawater.
Table 4. Formulations of Artificial Seawater (Added to 1 kg of Water)31 compound
weight (g)
NaCl MgCl2
23.939 5.079
Na2SO4 CaCl2 KCl
compound
weight (g)
KBr H3BO3
0.098 0.027
3.994
SrCl2
0.024
1.123
NaF
0.003
0.667
NaHCO3
0.196
open system and synthetic gas was bubbled in continuously, the CO2 partial pressure could be considered as 0.15 atm. The pH of the seawater was adjusted by adding ammonia/ammonium chloride buffer medium, which could resist dramatic changes in pH and maintain relatively constant alkaline conditions. The pH values were measured with a Mettler-Toledo pH meter. During the experiments, 10 mL of solution was withdrawn every 1 min by a micropump and placed in sealed bottles. After a rest period long enough for liquidsolid separation and crystallization (about 4 days), each suspension sample was filtered through 0.20-μm nucleopore polycarbonate membrane filters. The concentrations of Ca2þ and Mg2þ in the separated liquid phase were measured by ion chromatography. Several analysis methods were jointly applied to determine the micromorphologies, structures, compositions, and elemental contents of the precipitation products. X-ray powder diffraction (XRD) patterns were obtained on a diffractometer operated at 40 kV and 40 mA. The XRD patterns were recorded from 10° to 80° (2θ) at a rate of 0.06° per step. The micromorphology was investigated by scanning electron microscopy (SEM) using a JEOL JSM-7600F microscope operated at 0.530 kV and 1013(2 109) A that was equipped with an X-ray energy-dispersive spectroscopy (EDS) attachment. A Vario EI III
Figure 4. Kinetics of the carbonation reaction of gaseous CO2 with seawater.
element analyzer was employed to measure the contents of C/H/N in the precipitation products.
4. RESULTS AND DISCUSSION 4.1. Carbonation Process. Figure 4 reflects the reaction process specifically when CO2 was bubbled in the seawater. It can be noticed that, with the bubbling of CO2, the pH of the solution system declined slowly without any sudden changes, but there were obvious changes in the concentrations of Ca and Mg ions. The initial concentration of Ca2þ was about 420 mg/L, and there was a sharp decline at about the seventh minute. At the same time, the initial concentration of Mg2þ was about 1350 mg/L, and its sudden change emerged after about 60 min of bubbling of CO2. After 90 min, the concentrations of Ca and Mg ions remained nearly unchanged. Of course, if the bubbling of CO2 were continued, their concentrations might increase again because of the bicarbonation reactions. However, the carbonation 8336
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Industrial & Engineering Chemistry Research reaction was the focus here. The final concentrations of Ca2þ and Mg2þ were 12 and 120 mg/L, respectively, which were about 3% and 9% of their initial concentrations. This means that over 90% of the Ca and Mg ions in the seawater solution migrated into the precipitation products, which could be observed with flocculent shapes while the reactions were proceeding. The precipitation products should be calcium and/or magnesium carbonate. The carbonation process can be described as follows: In the beginning, CaCO3 and MgCO3 were unsaturated in the seawater solution because of the low concentration of CO32. With the bubbling of CO2, increasing numbers of CO32 ions were generated. Then, at the seventh minute, CaCO3 became supersaturated and began to precipitate, and with the continuous increase in the CO32 concentration, at about the 60th minute, MgCO3 began to precipitate. Finally, the increase of the CO32 concentration ceased, and the formation of precipitates stopped with only very small fractions of the Ca and Mg ions remaining in the solution. The initial pH of the solution was set to 10.25, because of the injection of ammonia/ammonium chloride buffer medium. The pH value was greater than the value of 9.73 calculated in section 2.4 at which Mg(OH)2 began to precipitate, but no obvious precipitation was observed in the primary stages of the experiment. The reason for this phenomenon is that the injection of NH3/ NH4Cl buffer increased the background electrolyte concentration of the seawater solution and accordingly lowered the activity coefficients of the ions, so the initial pH value for Mg(OH)2 to precipitate was raised. In addition, compounds under slightly supersaturated conditions might not precipitate from solutions because the reaction tendency might be very weak. In the experiments, it was found that ammonia/ammonium chloride buffer was better than strong alkali as the pH regulating agent. When NaOH was used to increase the pH, dramatic pH changes were encountered, and the experiments could not be conducted successfully with CO2 bubbling. It should be pointed out that it took nearly 100 min for the carbonation process to be finished, because of the low flow rate of gaseous CO2 and the relatively huge volume of seawater in the reactor. The purpose of the low-flow-rate design is to trace the reaction process and to understand changes more clearly. An experiment at higher CO2 flow rate, about 10 L/min, was also carried out. It was found that carbonate precipitation could be almost completed within about 10 min. Therefore, the reaction rate is acceptable for practical engineering implementation of this CO2 fixation idea. 4.2. Characterization of the Precipitation Product. During the experiments, the seawater was initially clear and transparent, but the emergence of flocculent precipitation gradually changed it into a suspension. The final suspension was set aside for about 4 days before being filtered, so that the precipitation products would be sufficiently crystallized. Then, the well-crystallized solid could be analyzed by XRD, SEM, EDS, and so on. Figure 5 shows the XRD diffractogram of the final precipitation sample. The narrow peaks indicate the high degree of crystallization. By comparison with JCPDS cards, it was determined that the peaks of CaMg(CO3)2 and MgCO3 3 nH2O predominated in the diffractogram. However, the discussion in section 4.1 indicates the precipitation of CaCO3 and MgCO3 occurred chronologically and that their time interval was about 50 min. Thus, step precipitation does not mean step crystallization; the CaCO3 and MgCO3 microparticles must have interacted with each other and generated the intergrowth crystal CaMg(CO3)2. Because the amount of MgCO3 was greater than the amount of CaCO3, the excess MgCO3 eventually existed in the form
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Figure 5. XRD diffractogram of the final precipitation product.
of MgCO3 3 nH2O. The MgCO3 usually contains three waters of crystallization, called nesquehonite (MgCO3 3 3H2O). However, the amount of crystallization water varied in our experiments, and accordingly, it was described as MgCO3 3 nH2O. SEM images of the solid sample are shown in Figure 6, where image a is magnified by a factor of 500 and image b is magnified by a factor of 400. Two typical crystal shapes were observed under the electron microscope: rod shapes and ball shapes. EDS analyses indicated that the rod-shaped particles should be MgCO3 and the ball-shaped particles should be CaMg(CO3)2. To further determine the compositions of the precipitation sample, elemental analyses were also performed, and Table 5 presents the results. Therein, the contents of Ca and Mg were obtained by titration, and those of C, H, and O were determined by elemental analysis. According to the initial and final concentrations of Ca2þ and Mg2þ of the seawater solution, the amounts that precipitated can be determined. The mass ratio of Ca to Mg should be 1:2.99. In Table 4, the ratio value was calculated as 1:3.01. Thus, it can be confirmed that all of the disappeared Ca and Mg ions entered the precipitate. In addition, the (Ca þ Mg)/C molar ratio was very close to 1:1 in the precipitated sample, which is agreement with the molar ratio of calcium and magnesium carbonates. Moreover, the total mass contents of elements Ca, Mg, C, H, and O amounted to 98.96%. This means that all of the precipitation products should be carbonates and that the margin of about 1% can be attributed to impurities in the sample or measurement errors. Therefore, all of the above analysis results confirm that calcium and magnesium carbonates were the main components of the precipitation sample.
5. PROSPECTS It can be concluded that the idea of CO2 fixation using Ca and Mg ions in solutions such as natural seawater is feasible in terms of both theory and kinetics under some specific conditions. This work highlights a permanent method for CO2 capture and sequestration. Based on the theoretical analyses and experiments reported herein, the potential of this method to fix CO2 can be evaluated. Because the decrease of the Ca2þ and Mg2þ concentrations is equal to the molar amount of carbon sequestered, eq 12 can be used to determine how much CO2 would be fixed per unit volume of seawater Cfixation ¼ ð½Ca2þ initial ½Ca2þ final Þ þ ð½Mg2þ initial ½Mg2þ final Þ 8337
ð12Þ
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Figure 6. SEM images of the final precipitation sample.
Table 5. Elemental Analysis Results for the Final Precipitation Sample element
content (mass %)
Ca
6.34
Mg
18.96
C H
11.33 1.89
O
60.44
total
98.96
The experiments showed that the migration ratios of Ca2þ and Mg2þ from the seawater solution to the precipitation products could exceed 90%. Assuming utilization of 90% of the Ca and Mg ions, the value of Cfixation can be calculated to be 0.06 mol of CO2 per liter of natural seawater. That is, the carbon fixation capacity of 1 m3 of natural seawater is about 1.34 m3 or 2.65 kg of gaseous CO2 at standard conditions. Moreover, the CO2 fixation potential of concentrated seawater will be 23 times this value. In 2009, the global CO2 emissions were 31.3 billion tonnes.32 If 10% of the total, or approximately 3 Gt, could be fixed by the Ca and Mg ions in natural seawater, the amount of seawater needed would be around 1130 km3, which is less than 1 ppm of the total seawater. This means that, even if the annual CO2 emissions of the entire world were all captured in this way, the concentrations of Ca2þ and Mg2þ in seawater would change at only the part-permillion scale. Therefore, the ecological effects might be negligible. At present, the desalination capacity worldwide is about 40 million m3 per day.33 The concentrated seawater generated from these desalination plants is also about 40 million m3 per day, most of which is presently poured back into oceans without utilization. If the idea of CO2 fixation is implemented, the reduced CO2 emissions would be about 100 Mt per year using concentrated seawater. Moreover, around 200 Mt of calcium and magnesium carbonates would also be produced as byproducts. This might provide a significant benefit in efforts to mitigate global climate change.
6. CONCLUSIONS Both theoretical analyses and experimental results confirm the possibility of our initial idea to fix CO2 using the Ca and Mg ions in enriched solutions such as seawater. The carbonate precipitation reactions can be enhanced by raising the pH or the CO2 partial pressure. In the experiments, over 90% of the Ca2þ and Mg2þ ions in the seawater could be converted to precipitation products in the forms of MgCO3 and dolomite [MgCa(CO3)2], and the whole process was found to be quite acceptable in terms
of kinetics. The implementation of this idea could realize not only the permanent fixation of CO2 but also the production of large amounts of carbonate byproducts. Although this novel idea seems sound, some technical points must be well researched and designed to realize practical application, such as the control of pH, the selection of buffering solutions, and the separation of the precipitates. If catalysts for the relevant reactions could be developed, the application of this technology would also be greatly advanced. Further work in these areas is just beginning in our laboratory. However, the aim of publication for this article was to encourage more efforts or collaborations to implement this technology as soon as possible in order to meet the global challenges posed by climate change.
’ AUTHOR INFORMATION Corresponding Author
*Tel.: þ86-531-88399372. Fax: þ86-531-88395877. E-mail:
[email protected].
’ ACKNOWLEDGMENT The authors acknowledge support from the National HighTech Research and Development Program of China (863 Program) (Grant 2009AA05Z303) and the Program for New Century Excellent Talents in University (NCET-10-0529). ’ REFERENCES (1) Chu, S. Carbon Capture and Sequestration. Science 2009, 325, 1599. (2) Kane, R. L.; Klein, D. E. Carbon sequestration: An option for mitigating global climate change. Chem. Eng. Prog. 2001, 97, 44–52. (3) Ishibashi, M.; Ota, H.; Akutsu, N.; Umeda, S.; Tajika, M.; Izumi, J.; Yasutake, A.; Kabata, T.; Kageyama, Y. Technology for Removing Carbon Dioxide From Power Plant Flue Gas by the Physical Adsorption Method. Energy Convers. Manage. l996, 37, 929–933. (4) Lyngfelt, A.; Leckner, B. Technologies for CO2 separation. Presented at the Minisymposium on Carbon Dioxide Capture and Storage, G€oteborg, Sweden, Oct 22, 1999. (5) Johansson, E.; Lyngfelt, A.; Mattisson, T. A circulating fluidized bed combustor system with inherent CO2 separation—Application of chemical looping combustion. In Proceedings of the 7th International Conference on Circulating Fluidized Beds; Grace, J. R., Zhu, J., de Lasa, H., Eds.; Canadian Society for Chemical Engineering: Ottawa, Canada, 2002; pp 717724. (6) AL-Saffar, H. B.; Ozturk, B.; Hughes, R. A comparison of porous and non-porous gasliquid membrane contactors for gas separation. Trans. Inst. Chem. Eng. A 1997, 75, 685–692. 8338
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’ NOTE ADDED AFTER ASAP PUBLICATION After this paper was published online June 6, 2011, a correction was made to a reference citation in Table 2. The corrected version was published online June 8, 2011.
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