Article pubs.acs.org/EF
Use of Reactive Species in Water for CO2 Mineralization Juan Ma and Roe-Hoan Yoon* Center for Advanced Separation Technologies, Virginia Polytechnic Institute and State University, 146 Holden Hall, Blacksburg, Virginia 24061, United States ABSTRACT: Various carbon capture and sequestration (CCS) technologies have been developed to address the issues concerning climate changes associated with anthropogenic CO2 emissions. In the present work, possibilities of mineralizing CO2 with the reactive species, such as Mg2+ ions, present in nature, such as seawater and produced water, have been explored. Laboratory tests conducted with solutions containing 1400 ppm of Mg2+ ions showed that nesquehonite [Mg(OH)(HCO3)·2H2O] is formed upon CO2 injection to the solution at an atmospheric pressure. The results showed that, for mineralization to occur, the pH should be raised above 6.8, as predicted from thermodynamics. Kinetic studies conducted at different temperatures showed that the nesquehonite formation involves an activation energy of 66.7 kJ/mol, which can be overcome by increasing the mass- and heat-transfer efficiency as well as the operative temperature. On the basis of the kinetics data obtained at a low agitation speed, the number and volume of the mineralization reactors required to capture CO2 emitted from a 600 MW coal power plant have been determined. In addition, the amount of alkalis needed to raise the pH for precipitation and, subsequently, to obtain the natural pH of seawater has been estimated. CCS cause the electricity costs to increase by ∼76%, which may decrease in time to ∼35% via “learning-by-doing”. Many investigators explored the possibility of reacting CO2 with basic minerals and storing the resulting carbonates (or minerals). Because carbonates are stable on a geological time scale, they can be stored permanently under ambient conditions with minimal precautions. Ultramafic rocks, such as basalt and peridotite, may be used for the mineralization. The major components of these rocks are olivine (Mg2SiO4), pyroxene (MgSi2O6), and calcium feldspar, of which magnesium (Mg), iron (Fe), and calcium (Ca) are the most important reactive species that can form carbonates. Basalt is the common extrusive igneous rock formed on the crustal surface; however, its reaction rate is slow because of its fine-grained texture. It has been shown that the reaction rate of coarse-grained peridotite is fast enough for CO2 storage by in situ mineralization. In this approach, supercritical CO2 is injected through bore holes drilled into the intrusive igneous rocks.6 The reaction rate can be accelerated once the carbonization, which is exothermic, is initiated and the temperature reaches 150−200 °C. Lackner et al.7,8 developed an ex situ process of mineralizing CO2, in which magnesium hydroxide [Mg(OH)2] is extracted from olivine by heat and acid (HCl) treatments and the hydroxide is contacted with CO2 gas at a high temperature (140−300 °C) and pressure to form magnesium carbonates. However, the process is energy-intensive, generating 4 times as much CO2 as that to be captured.9,10 O’Connor et al.11−13 developed a less costly process, in which ultramafic rocks are “activated” such that reactive species, such as Mg2+, Ca2+, and Fe2+ species, are more readily released into aqueous media.
1. INTRODUCTION According to the Fourth Intergovernmental Panel on Climate Change (IPCC) Assessment Report (AR4),1 global warming is unequivocal, as evidenced by increasing air and ocean temperatures, melting of artic snow and ice, rising sea levels, and frequent extreme weather events. The major culprit of these problems is considered to be the CO2 emitted from burning fossil fuels. The report suggests four abatement options, which include fuel switch, energy efficiency, renewable and nuclear energies, and CO2 capture and sequestration (CCS). The last is essential for continued use of fossil energy for electricity and heat generation. At present, fossil energy accounts for 67% of electricity generated globally and 90% of the 33.9 gigatonnes of CO2 emitted in 2011.2 In the CCS option, CO2 is separated from flue gas and compressed to a liquid state (capture), transported via a pipeline (transport), and pumped underground for storage (storage). Engineering studies suggest that the capture portion of the process will account for approximately 90% of the additional costs associated with CCS.3 The storage sites include oil and gas wells, deep saline aquifers, and unminable coal seams. Dependent upon the nature of the storage site, CO2 is stored as compressed gas, liquid, or supercritical fluid. Supercritical CO2 is immiscible with water, easy to pump because of low viscosity, and has a high affinity for hydrophobic substances, such as coal and oil, a property that is important for the enhanced oil recovery (EOR) and methane (CH4) recovery from unminable coal seams. The four different geological formations noted above have an estimated capacity of ∼2000 gigatonnes of CO2 globally. In addition, CO2 can be stored in deep oceans, where it becomes heavier than water and stays at the bottom. Estimated costs for the capture, transport, and storage are $20−95, $1−10, and $0.5−10/tonne of CO2 without including long-term monitoring costs, respectively.4,5 The additional costs associated with © 2013 American Chemical Society
Special Issue: Accelerating Fossil Energy Technology Development through Integrated Computation and Experiment Received: February 1, 2013 Revised: April 15, 2013 Published: April 23, 2013 4190
dx.doi.org/10.1021/ef400201a | Energy Fuels 2013, 27, 4190−4198
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Table 1. Standard Gibbs Free Energy of Formation (ΔGf°) at 298 K Used for the Construction of Solubility Diagrams and Thermodynamic Analysis composition
ΔGf° (kJ mol−1)
reference
composition
ΔGf° (kJ mol−1)
reference
CO2(g) H2O(l) OH− H2CO3 HCO3− CO32− Mg(OH)(HCO3)·2H2O(s) Mg2+ MgOH+ Mg(OH)2(aq) Mg(OH)2(s) MgCO3(aq) Mg2CO32+a MgHCO3+
−394.39 −237.14 −157.2 −623.1 −586.85 −527.8 −1723.95 −453.99 −624.49 −769.4 −833.7 −1000.3 −1456.27 −1048.35
27 27 27 27 27 27 31 29 29 27 27 32
Mg2SiO4(s) MgHSiO3+a SiO2(aq) SiO2(s), (amorphous) H4SiO4 H3SiO4− H2SiO42− CaCO3(s) Ca2+ CaCO3(aq) CaHCO3+ CaOH+ Ca(OH)2(s) Ca(OH)2(aq)a
−2056.5 −1476.99 −833.41 −850.56 −1309.18 −1253.01 −1176.64 −1129.1 −552.79 −1099.13 −1145.99 −716.62 −897.5 −883.51
28
32
29 30 30 29 29 31 29 31 31 29 27
ΔGf° values of Mg2CO3 , MgHSiO3 , and Ca(OH)2(aq) are calculated using the following reactions a, b, and c, respectively: (a) 2Mg2+ + CO32− = Mg2CO32+, pKa = −3.59; (b) Mg2+ + SiO2(aq) + H2O = MgHSiO3+ + H+, pKb = 8.33; and (c) Ca(OH)2(aq) = Ca(OH)2(s), pKc = −2.45, where Ka is from the computer program Visual Minteq, Kb is from the study by Giammar et al.,33 and Kc is from the study by Somasundaran and Agar.34 a
2+
+
that can be used to mix pure CO2 gas into MgCl2 solutions to form nesquehonite [Mg(OH)(HCO3)·2H2O] under ambient conditions. The authors used aqueous ammonia (NH3) to control the pH in the range of 7.8−8.2. Approximately 95% of the magnesium ions were precipitated as nesquehonite within 10 min in a 7 g/L Mg solution. It has been shown also that brine solutions containing Ca2+ and Mg2+ ions can be used to mineralize CO2 using sodium hydroxide (NaOH) produced using a proprietary electrochemical process as an alkali.25,26 In the present work, kinetics of carbonation of Mg2+ ions in aqueous media has been studied. Because the ocean is the largest source of the reactive species, experimental work has been carried out at the average Mg2+ ion concentration (1400 mg/L) in seawater. The kinetic studies have been carried out at different temperatures to determine the activation energy of carbonation of Mg2+ ions, which is useful for understanding the mechanisms involved and for estimating the size and number of reactors required for mineralizing the CO2 emitted from a 600 MW coal-burning power plant using seawater. Thermodynamic analysis has also been carried out to determine the critical pH of precipitation and estimate the amount of alkalis required for mineralization. The fundamental information obtained in the present work may also be useful for better understanding the mechanisms involved in the carbonation of saline water.
They are carbonated by reacting with dissolved CO2 at elevated pressures (40−150 atm) and temperatures (100−180 °C). The activation steps include fine grinding of basalt rock, acid treatment, and heat treatments to facilitate the release of the reactive species. The estimated costs of the wet carbonation process are in the range of $50−100/tonne of CO2, which translate into $10−30 energy penalty on the original power plant. When a 10−40% energy penalty is taken into account a full CCS system with the wet carbonation process would require 60−180% more energy than a power plant with equivalent output without CCS. Further, the process requires 5 times as much rocks as coal, which will severely limit its applicability.14 Other investigators15−18 developed similar processes to extract the reactive species into solution using various acids, bases, and chelating agents, with mixed results. It appears that the major difficulty associated with CO2 mineralization is the high cost of extracting reactive species, which is not surprising in that the ultramafic rocks are thermodynamically stable. One way to circumvent this problem would be to identify appropriate sources of reactive species that can be carbonated without activation. These include seawater containing ∼1400 mg/L Mg2+ ions and ∼400 mg/L Ca2+ ions, wastewater from desalination plants containing typically 3000 mg/L Mg, produced water from oil and gas industries, saline water containing high concentrations of divalent cations, etc. The brine water from the Oriskany Sandstone aquifer in Indiana and Pennsylvania, for example, contains 2490 mg/L Mg2+ ions.19 Another example is the water from a 1158 m deep well in Guernsey, OH, containing 3440 mg/L Mg2+ ions.20 The ocean is the largest CO2 sink, with most of the gas dissolved from the atmosphere staying in the surface water. If CO2 gas is injected into deep oceans, approximately 103−104 gigatonnes of carbon can be dissolved, while the current anthropogenic emissions are in the range of 6−8 gigatonnes of carbon per year.21 When carbon is emitted to the atmosphere, about 80% is transferred to the oceans on a time frame of ∼300 years, with the remainder being absorbed much more slowly. It has been suggested that the process can be accelerated by raising the pH to allow for precipitation of MgCO3 and CaCO3.22 At pH above 9.25, magnesium precipitates as hydroxide. Ferrini et al.23,24 designed a laboratory-scale reactor
2. THEORETICAL BASIS 2.1. Thermodynamics. Carbonation of Mg2+ ions as studied in the present work involves two steps. In step 1, CO2 is dissolved in a magnesium chloride (MgCl2) solution as follows: CO2 + H 2O → H+ + HCO3−
(1) −7.83
whose equilibrium constant (K1) is 10 at 25 °C, as obtained from the thermodynamic data presented in Table 1. In step 2, pH is raised above the critical pH of precipitation, so that the bicarbonate ions formed in reaction 1 combine with the Mg2+ ions in solution to form nesquehonite Mg 2 + + 2HCO3− + 2H 2O → Mg(OH)(HCO3) ·2H 2O(s) + CO2 4191
(2)
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with K2 = 103.24 at 25 °C. As will be shown later, the precipitates formed under the experimental conditions employed in the present work were indeed nesquehonite. From reactions 1 and 2, one obtains the overall carbonation reaction as follows:
as the critical pH of precipitation. Thus, nesquehonite should precipitate when CO2 gas is blown into a solution containing 1400 mg/L of Mg2+ ions while maintaining the pH above 6.8. 2.2. Kinetics. Assuming that nesquehonite formation (reaction 2) is a first-order process, its kinetics may be represented by the following relation:
Mg 2 + + CO2 + 4H 2O → Mg(OH)(HCO3) ·2H 2O(s) + 2H+
dC = k(C0 − C) dt
(3)
According to the thermodynamic data given in Table 1, K3 = 10−12.41. Reaction 1 is fast, while reaction 2 is slow. Therefore, the objective of the present investigation was to study the kinetics of reaction 2. For the nesquehonite formation to be spontaneous, the concentrations of Mg2+ and HCO3− ions should be such that its reaction quotient (Q) is less than K2 = 103.24, i.e., Q=
pCO2 2+
[Mg ][HCO3−]2
where C0 is the initial concentration of the Mg ions in a reactor, C is the concentration at time t, and k is the rate constant. When eq 8 is integrated, one obtains
⎛C⎞ ln⎜ ⎟ = −kt ⎝ C0 ⎠
< 103.24 (4)
⎛ E ⎞ k = A exp⎜ − a ⎟ ⎝ RT ⎠
In the present work, [Mg ] = 0.058 mol/L and pCO2 = 1 atm. When these values are substituted into eq 4, one finds that nesquehonite should form spontaneously when [HCO3−] > 0.1 mol/L. Reaction 1 is thermodynamically equivalent to the following reaction: CO2 + OH →
HCO3−
(9)
When ln(C/C0) versus t is plotted, one can determine the value of k by monitoring C as a function of time in the experiment. According to the Arrhenius equation
2+
−
(8) 2+
(10)
where k is the rate constant, A is a constant, Ea is the activation energy, R is the gas constant, and T is the absolute temperature. From eq 10, one can derive the following relation: E d ln k =− a d(1/T ) R
(5)
which shows that the concentration of the bicarbonate ions should increase with pH. Figure 1 shows the changes in the
(11)
that can be used to determine Ea from a ln k versus 1/T plot, which is useful for understanding the mechanism(s) of nesquehonite precipitation and for predicting k at different temperatures.
3. EXPERIMENTAL SECTION 3.1. Materials. Anhydrous MgCl2 (99%, Alfa Aesar) was used to prepare aqueous solutions containing Mg2+ ions, while a compressed CO2 gas (bone-dry grade, Airgas) was used as a source of carbon dioxide. The solution pH was controlled by adding aliquots of NaOH solutions. Diethylenediaminetetraacetic acid disodium salt (Na2H2EDTA·2H2O, >99%, Alfa Aesar), pure Eriochrome Black T (EBT, indicator grade, Acros Organics), 2-methoxyethanol (ACS, >99.3%, Alfa Aesar), and a pH 10 ammonia buffer solution (SigmaAldrich) were used to determine the concentration of Mg2+ ions in solution. All solutions were prepared using the Millipore ultrapure water with a resistivity of 18.2 MΩ/cm, which was obtained using a Direct-Q3 water purification system. 3.2. Procedure. In step 1, a stream of CO2 gas was injected into a MgCl2 solution (700 mL, 1400 mg/L) under ambient conditions, i.e., at room temperature and 1 atm, while being agitated by means of a magnetic stirrer. The natural pH of the MgCl2 solution was ∼8.3, which dropped precipitously to pH ∼5 upon CO2 injections. The pH of the solution was raised to a desired value by adding aliquots of NaOH solutions, while the CO2 injection and agitation continued. Once a desired pH was reached, both the CO2 injection and agitation were stopped and the solution was left to stand for 10 min to allow for the system to reach an equilibrium. During this time, no precipitates were formed and the pH did not change. In step 2, the solution containing the HCO3− ions formed in step 1 via reaction 1 and the Mg2+ ions originally present in solution were agitated to facilitate the formation of nesquehonite via reaction 2. Samples of the solution were taken at different reaction times, filtered, and analyzed for residual Mg2+ ion concentration in solution using the EDTA titration method.35 The precipitates were characterized by means of X-ray diffraction (XRD) and field-emission scanning electron microscopy (FE-SEM) analyses. Figure 2 shows the XRD spectra and SEM images obtained on the precipitates obtained at three different
Figure 1. Effect of pH on the concentration of carbonate species in water at 25 °C and pCO2 = 1 atm. The dotted line represents the total carbonate concentration.
concentration of the HCO3−, CO32−, and H2CO3(aq) species with pH under the experimental conditions carried in the present work, i.e., pCO2 = 1 atm. The dotted line represents the total carbonate concentration. Reaction 1 gives the following relation: K1 =
[H+][HCO3−] pCO2
(6)
which, in turn, gives the following equation: pH = log[HCO3−] − log pCO2 − log K1
(7)
that can be used to calculate the critical pH of nesquehonite precipitation at different concentrations of HCO3− and CO2 partial pressures. When the values of [HCO3−] > 0.1 mol/L, pCO2 = 1, and K1 = 10−7.83 are substituted, one obtains pH 6.8 4192
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Figure 2. XRD patterns (left) and SEM images (right) of solid products obtained at different temperatures of (a) 20 °C, (b) 30 °C, and (c) 40 °C. Experiments were conducted at a stirring speed of 640 rpm, an aeration rate of 10 L/min, and initial step 2 pH of 6.9. temperatures, i.e., 20, 30, and 40 °C. The initial pH was 6.9, and the aeration rate was kept constant at 10 L/min. The X-ray spectra of the samples correspond to that of nesquehonite (ICSD 01-070-6308), and the precipitates were needle-like. At 40 °C, however, smaller plate-like crystals were also formed possibly because of the faster kinetics of precipitation.
Part of the hydroxyl ions formed in this manner may also be used to precipitate nesquehonite. 4.2. Effect of Agitation. Figure 4 shows the effect of the stirring speed on the changes in Mg2+ ion concentration during
4. RESULTS AND DISCUSSION 4.1. Effect of the Reaction Time. Figure 3 shows the changes in the Mg2+ ion concentration, pH, and weight
Figure 4. Effect of the agitation speed on the kinetics of nesquehonite precipitation. The experiments were conducted at 40 °C and the initial pH of 6.9. The kinetics increases with agitation speed. Figure 3. Effect of agitating a 1400 mg/L Mg solution at 640 rpm on changes in Mg2+ ion concentration, percent utilization of Mg2+ ions, and pH. The experiment was conducted at 40 °C, and the initial pH was 6.9.
step 2. As shown, the concentration decreased faster at higher speeds. At 640 revolutions per minute (rpm), it took 60 min to reduce the concentration from 1400 to 260 mg/L, while it took only 20 min to reach approximately the same concentration at 1300 rpm. Thus, the rate of nesquehonite precipitation, as represented by the changes in the Mg2+ ion concentration, increased substantially with an increasing stirring speed, most probably because of the increased mass-transfer rates of the reactants and products involved. As shown by reaction 2, CO2 evolves during nesquehonite formation. One way to increase the kinetics would be to dilute the gaseous product by blowing air into the solution. Figure 5 shows the results obtained by injecting air at 10 L/min at 640 rpm and 40 °C. As shown, the kinetics of nesquehonite formation increased considerably by blowing air. The aeration may have not only helped decrease the CO2 concentration in solution but also helped increase mixing, both of which should have contributed to increased kinetics of nesquehonite precipitation. 4.3. Effect of NaCl. Recognizing that seawater is probably the largest source of Mg2+(aq) species on earth, a series of kinetics tests was conducted in the presence of 0.43 mol/L NaCl, with the results presented in Figure 6. As shown, the salt caused a decrease in the percent utilization of Mg2+ ions. One possible reason for this may be that, in the presence of NaCl,
percentage (wt %) of the Mg2+ ions carbonated during step 2. In this particular experiment, the pH at the beginning of step 2 was 6.9, which increased as nesquehonite began to precipitate. As the precipitation continued, the Mg concentration decreased and pH increased with stirring time. After 1 h of agitation, the Mg2+ ion concentration was reduced from the initial 1400 to ∼200 mg/L, which represented 86% utilization of Mg2+ ions. Note that the pH increased from 6.9 to ∼9.3 during step 2, which may be attributed to the dissociation of part of nesquehonite formed via reaction 2. As will be discussed later in conjunction with its solubility diagram (Figure 12), nesquehonite can serve effectively as a pH 9.3 buffer. It can be readily shown that the amount of nesquehonite to be dissociated to raise the pH from 6.9 to 9.3 is only a small fraction (
