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Energy & Fuels 2006, 20, 172-179
Investigation of Carbon Sequestration via Induced Calcite Formation in Natural Gas Well Brine Matthew L. Druckenmiller,*,† M. Mercedes Maroto-Valer,‡ and Meredith Hill† The Energy Institute, The PennsylVania State UniVersity, UniVersity Park, PennsylVania 16802, and School of Chemical, EnVironmental and Mining Engineering, UniVersity of Nottingham, UniVersity Park, Nottingham, NG7 2RD, U.K. ReceiVed April 21, 2005. ReVised Manuscript ReceiVed October 17, 2005
The permanent storage of carbon in mineral form using natural brines found in geologic formations is at the forefront of carbon sequestration research. A complex chemistry describes the ultimate fixation of carbon in stable minerals, such as calcite. However, the parameters that govern carbonate formation are not well understood. Accordingly, the purpose of this study is to induce and characterize calcite formation by reacting natural gas brine with CO2. Brine pH has a significant effect on this conversion and can thus be adjusted to induce calcite precipitation using a laboratory scale reactor operated at temperatures of 75 and 150 °C and pressures of 600 and 1500 psi. Initial pH conditions of at least 9.0 are optimal for carbonate precipitation in reactions of 18 h. Although the reaction duration is not long enough to successfully correlate brine compositional changes with precipitation and pH evolution, X-ray diffraction analysis clearly confirms the presence of calcite. Scanning electron microscopy/energy-dispersive X-ray spectroscopy analysis provides an introductory look at the microscale production of these minerals.
Introduction In 2001, United States CO2 emissions accounted for 1.57 GtC, which was 24% of the total global carbon releases from CO2.1 Of that 1.57 GtC, 36% was from energy- and industry-related coal combustion. It is also projected that by the year 2025 coal consumption by U.S. electricity generators will increase by 48% from the 2001 level.1 Carbon sequestration, along with increases in efficiency and the advancement of clean-coal technologies, may serve to minimize the environmental impact of coal. With the proximity of U.S. electricity generators to deep saline aquifers of large estimated capacity, geologic sequestration is proposed as a suitable method to reduce CO2 emissions. Capability exists in the technologies currently being explored by energy industries regarding concentrating and transporting CO2 from flue gas, fluid injection into subsurface media, and the characterization of geologic formations. U.S. deep saline aquifers are estimated to provide storage for approximately 130 GtC equivalent, which is approximately 80 times the United States’ total carbon emissions in 2001.1,2 Geologic sequestration in saline aquifers is a complex process with multiple mechanisms for carbon storage. Saline aquifers are deep porous rock formations that are saturated with brine, which is typically rich in various metals. Following the injection of CO2 into the subsurface below a typical depth of 800 m, the mechanism for CO2 storage is initially hydrodynamic as the CO2 is stored as a dense supercritical fluid. Following the dissolution of CO2 into the liquid phase, chemical interactions * Corresponding author. Address: P.O. Box 80054, Fairbanks, AK 99708. E-mail:
[email protected]. † The Pennsylvania State University. ‡ University of Nottingham. (1) Energy Information Administration, DOE. International Energy Annual 2002. http://www.eia.doe.gov/emeu/iea/carbon.html (accessed March 30, 2005). (2) Carbon Sequestration - Research and DeVelopment; DOE/SC/FE-1; U.S. Department of Energy: Washington DC, 1999.
with dissolved metals may form stable mineral carbonates on a much longer, yet currently unknown, time scale.3 CO2 may react with brine’s metal cations, such as calcium and magnesium, to form carbonate mineral precipitates. The following shows a simplified reaction sequence in which calcite and magnesite are ultimately formed in reactions 5a and 5b, respectively:4,5
CO2(g) S CO2(aq)
(1)
CO2(aq) + H2O S H2CO3
(2)
H2CO3 S H+ + HCO3-
(3)
HCO3- S H+ + CO32-
(4)
Ca2+ + CO32- S CaCO3 V
(5a)
Mg2+ + CO32- S MgCO3 V
(5b)
Reaction 1 is the dissolution of CO2 in water, which is highly dependent on temperature, pressure, and brine salinity.5 Water may then react with CO2 to form carbonic acid as shown in reaction 2. Reaction 3 represents the dissociation of carbonic acid and the liberation of protons, which reduces the pH of the system. The bicarbonate ion may then dissociate via reaction 4 to form the carbonate ion. The metals then react with the carbonate ion to form the carbonate minerals.5 The pH of the system has an extensive influence on the solution chemistry. The proportions of the carbonic species are controlled by pH as it determines the dominant step of the (3) Bachu, S. Energy ConVers. Manage. 2000, 41, 953-970. (4) Bond, G. M.; McPherson, B.; Stringer, J.; Wellman, T.; Abel, A.; Medina, M. Prepr. Pap.-Am. Chem. Soc., DiV. Fuel Chem. 2002, 47, 39. (5) Soong, Y.; Jones, J. R.; Hedges, S. W.; Harrison, D. K.; Knoer, J. P.; Baltrus, J. P.; Thompson, R. L. Prepr. Pap.-Am. Chem. Soc., DiV. Fuel Chem. 2002, 47, 43.
10.1021/ef050115u CCC: $33.50 © 2006 American Chemical Society Published on Web 11/25/2005
Carbon Sequestration Via Calcite Formation in Brine
reaction sequence.4 At a low pH (∼4) the production of H2CO3 dominates, at a mid pH (∼6) HCO3- production dominates, and at a high pH (∼9) CO32- dominates.5 Accordingly, a high pH favors the precipitation of carbonate minerals due to the availability of carbonate ions. The pH also determines the ratelimiting step of the reaction sequence. In the low-to-mid pH range, it is the hydration of CO2 to form carbonic acid in reaction 2, which has a forward reaction rate constant of 6.2 × 10-2 s-1 at 25 °C.6 However, in the high pH range the rate-limiting step is the first dissociation of carbonic acid to form bicarbonate as in reaction 3.5 Soong et al. suggest that pH has a significant effect on not only conversion rates but also the species of the precipitates.7 The culmination of the complex chemistry is that the rate of the mineral trapping process is slow and serves as the major disadvantage of this technology.5 However, it is known that the formation of carbonates can be promoted by increasing brine pH through the addition of a strong base.5,7 Although saline aquifers have the essential characteristics of a long-term solution, little is known about the kinetics of trapping CO2 in stable minerals.5 The parameters that affect the rate of the carbonate forming process are brine composition, temperature, pressure, and pH.5,6,8 For example, high temperature promotes the formation of carbonates because the dissociation of carbonic acid decreases with increasing temperature, thus increasing the pH.9 These parameters must be further understood since they determine the economical applicability of the technology as well as help to identify geologic locations that favor sequestration. We have previously examined the effects of temperature, pressure, and pH on the formation of carbonates during the reaction of CO2 with natural gas well brine using a laboratoryscale high-pressure/high-temperature reactor.8 A series of 16 6-h experiments were conducted at pressures of 600 and 1500 psi, temperatures of 75 and 150 °C, and initial brine pH of approximately 4.8 to 9.0. It was found that temperature had a greater influence on the evolution of the system’s pH throughout the course of the reactions as opposed to pressure. Therefore, it was concluded that temperature has a greater influence on carbonate mineral formation than does pressure, since pH has been identified as playing a major role in the formation of carbonates.8 Final brine liquid products and brine samples extracted during the reactions were analyzed using inductively coupled plasmaatomic emission spectroscopy (ICP-AES) analysis. The results of that analysis to correlate changes in brine composition with pH changes and the precipitation of carbonates were inconclusive as it was suggested that a longer reaction time was needed. Although the presence of carbonate mineral products, specifically calcite, was verified using X-ray diffraction (XRD) analysis in our previous work, no detailed characterization of the reaction’s solid products was conducted. Accordingly, the purpose of this study was to induce and characterize the natural geologic processes of mineral carbonate formation in the laboratory by reacting a natural gas well brine with CO2. It should be noted that this laboratory process is simplified from that of a natural in situ system due to the absence of a formation (6) Bond, G. M.; Stringer, J.; Brandvold, D. K.; Simsek, F. A.; Medina, M.; Egeland, G. Energy Fuels 2001, 15, 309-316. (7) Soong, Y.; Goodman, A. L.; Hedges, S. W.; Jones, J. R.; Harrison, D. K.; Zhu, C. 19th Annual International Pittsburgh Coal Conference, Pittsburgh, PA, Sept 23-27, 2002. (8) Druckenmiller, M. L.; Maroto-Valer, M. M. Carbon sequestration using brine of adjusted pH to form mineral carbonates. Fuel Process. Technol. 2005, 86, 1599-1614. (9) Read, A. J. J. Solution Chem. 1975, 4, 53-70.
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rock, which would present multiple additional reactions to consider by introducing various other minerals. However, such a simplified system is highly relevant to an ex situ industrial process that will lack a complex mineral matrix. After a reaction period of 18 h under controlled conditions of pH, temperature, and pressure, the formation of carbonates was investigated through changes in brine solution composition and the characterization of precipitates. The findings and experimental methods from our previous work were expanded upon and further investigated in this work to better understand the proposed correlations.8 The methods and results of the continued investigation are presented here. Experimental Section After a previous examination of the ability of various brine samples to maintain a raised pH value following treatment with KOH prior to reaction with CO2, we concluded that a specific brine sample, referred to as OH-1, was the most stable of the investigated brines.8 Results suggested that a relatively low iron composition in the liquid phase was the cause.8 This is because the oxidative precipitation of iron oxide, which is the transition of the soluble divalent ferric ion Fe2+ to the oxidized insoluble trivalent form Fe(III) oxide, is responsible for reducing the pH in the brines with a relatively high iron composition.8,9 Accordingly, OH-1 was selected as the experimental brine sample to be used in the extended CO2/ brine reactions discussed herein. The OH-1 brine sample was collected from a 1158-m natural gas well in Guernsey, OH, and supplied by the U.S. DOE National Energy Technology Laboratory.10 This section discusses the extended reactions between CO2 and brine, the methods used to characterize both unreacted brine and liquid brine by ICP-AES, and the methods used to characterize the solid products by XRD and scanning electron microscopy/ energy-dispersive X-ray spectroscopy (SEM/EDS). High-Pressure/High-Temperature Reactions between CO2 and Brine. A 180-mL Parr reactor (model series 4576, T316 stainless steel, custom 1.5 in. internal diameter) was used to react CO2 of 99.99% purity with the natural gas well brine OH-1 at high pressure and temperature. Figure 1 shows a schematic of the reactor system. Inspection of the reactor following trial experiments and throughout the course of the study suggested that the reactor is not prone to corrosion from the experimental conditions. As an extension of our study in which 6-h CO2/brine reactions were conducted, the experimental reaction time was extended to 18 h.8 Four experiments, labeled A to D, were conducted at pressures of 600 and 1500 psi, a temperature of 150 °C, and initial brine pH ranging from approximately 4.9 to 9.0. To promote adequate interaction, the reactor was constantly stirred at 400 rpm. In each experiment, liquid sampling occurred via the liquid sampling valve at regular time intervals to provide a snapshot of the brine pH and composition throughout the experiments. Brine samples of 95 mL were treated with KOH to the desired pH and immediately placed into the bomb cylinder and sealed. After placing the bomb cylinder in the heater assembly, the thermocouple, stirrer drive system, and water coolant supply were connected. The system was purged four times with CO2 at 600 psi. The system was then heated to the desired temperature. Next, the system was pressurized. For the experiments conducted at 600 psi, total pressure was supplied by CO2. For the experiment performed at 1500 psi, 600 psi of CO2 was initially charged, with the remaining pressure requirement supplied by N2 of 99.999% purity. Although a specific and finite volume of CO2 was charged to the reactor at the beginning of each experiment and following each sampling, it can be assumed, given the rates of carbonate formation, that the system operated as an open system with a constant supply of CO2. System temperature and pressure were monitored and recorded approximately every hour. A total of six samples of approximately (10) Soong, Y.; Fauth, D. J. National Energy Technology Laboratory, U.S. DOE, Pittsburgh, PA. Personal communication, 2003.
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Figure 1. Schematic of the high-pressure/high-temperature reactor. Reprinted with permission from ref 8. Copyright 2005 Elsevier.
10 mL were extracted during the course of the reactions. To minimize cooling and degassing, the pH of the extracted samples was immediately measured. The samples were then placed under refrigeration prior to further analysis. The solid products and the final liquid sample, however, were collected at the end of each experiment following depressurization and opening of the bomb cylinder. Once again, the pH of the solution was immediately measured. Next, the brine underwent gravity filtration to remove the solid precipitates. Following drying, the filtered solids were ground to a fine powder in preparation for characterization. Last, a sample of the filtrate was collected for ICP-AES analysis. Characterization of Unreacted Brine and Liquid Brine Products. A Leeman Labs PS3000UV ICP spectrometer was used to perform ICP-AES analysis on the brine samples. To establish an experimental baseline for brine composition, the unreacted brine underwent screening for various metals. The screening without acid pretreatment determined the elemental concentrations in the liquid phase, while screening with acid pretreatment to dissolve suspended solids using HNO3 provided total concentrations in solution.8 A quantitative analysis of the reacted brine samples using ICPAES screening was done to reveal how the elemental composition of brine changed during reaction with CO2. ICP-AES without acid pretreatment was used to analyze the 10-mL liquid samples extracted during the 18-hour reactions and those collected at the end. Prior to analysis, all samples were gravity-filtered using a filter with a 16-µm particle retention size and stored in a refrigerator at 5 °C. Characterization of Solid Products Using XRD. XRD is a nondestructive solid analysis method that employs the characteristic bending of X-rays to identify crystalline materials and thus was used to identify the minerals precipitated during the reactions. Solids recovered from the reactor experiments were characterized using a Scintag X-ray diffractor, which used wavelengths of 1.5406 Å. Analysis was conducted over an angle range of 20-60° and at a scanning rate of 1.20°/min. Very small amounts of sample (∼50 mg) were necessary for analysis as only enough was needed to form a thin film of solids on a quartz slide. SEM/EDS Analysis. Solids recovered from a reactor experiment were also viewed using a Hitachi S-3500N scanning electron microscope, which provided a magnified image of the solid surface. Due to the low conductivity of the precipitates, the specific samples analyzed were sputtercoated with gold to increase conductivity and to allow for a clearer image. EDS analysis, which identified and quantified the elemental composition of an area on the solid surface, was used in conjunction with SEM. The lateral resolution of the spectrometer was 1 µm.
Table 1. Experiments Conducted To React CO2 with Brine in the High-Pressure/High-Temperature Reactora experiment
pressure (psi)b
temp (°C)
initial pHc
final pH
A B C D
1500 600 600 600
150 150 150 150
9.02 8.97 7.11 4.85
5.41 5.08 5.51 5.06
a No liquid sampling occurred. Final liquid samples and precipitates were collected. b Pressures of 1500 psi represent a CO2 partial pressure of 600 psi and the remaining pressure supplied by N2. c pH measurements were taken of the brine following refrigeration at 5 °C.
Results and Discussion High-Pressure/High-Temperature Reactions between CO2 and Brine. Table 1 summarizes the four 18-hour experiments, labeled A through D, conducted in the reactor. No trend appears to exist between the starting and final brine pH values. Experiments B, C, and D were all conducted at 600 psi and 150 °C. Experiment C with a midrange initial pH of 7.11 had the highest final pH of 5.51, while the experiments with the highest and lowest initial pH values, experiments B and D, respectively, both had lower final pH values of approximately 5.1. The final pH values reported are meant for a qualitative comparison between the experiments. Because the brine samples immediately began to degas and cool following removal from the reactor, the measured pH, although done as quickly as possible, does not represent the actual in situ pH under the elevated temperature and pressure conditions. Any degassing of dissolved CO2 will result in an increase in pH. Future work with this reaction system aims to address this issue by using PHREEQC, an aqueous geochemistry modeling program that will employ speciation, solubility, and inverse modeling calculations to determine the in situ pH at specified temperatures and pressures. Although recent research11 has addressed the significance of elevated temperatures and pressures on the solubility and speciation of CO2 after injection, the authors fail to account for the evolution of solution pH after CO2 injection. Using a customized Pitzer database, PHREEQC can be used for such a purpose, and its use is currently being investigated by the authors.12 However, while the Pitzer database available now is more reliable than the generic databases that come with the code, (11) Allen, D. E.; Strazisar, B. R.; Soong, Y.; Hedges, S. W. Fuel Process. Technol. 2005, 86, 1569-1580.
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Figure 2. Concentration of various metals and pH versus time for experiments B through D. Segment 1 (not drawn to scale) represents the purging of the system four times with 600 psi of CO2, which lasted 10 min. Segment 2 (drawn to scale) represents the heating of the system from 25 to 150 °C, which took 30 min. Upon obtaining a temperature of 150 °C, the system was pressurized to 600 psi. Table 2. Characterization Data for the Experimental Brine Sample OH-1a
Ba Ca Fe K Mg Na Sr
metal concentration (ppm) without acid pretreatment
metal concentration (ppm) with HNO3 pretreatmentb
