172
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.
Energy & Fuels, Vol. 20, No. 1, 2006 173
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.
174 Energy & Fuels, Vol. 20, No. 1, 2006
Druckenmiller et al.
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.
Carbon Sequestration Via Calcite Formation in Brine
Energy & Fuels, Vol. 20, No. 1, 2006 175
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
<5
19570 ((500)
9 ((10)
2225 ((60)
3440 ((500)
69660 ((870)
2000 ((38)
5.66
22500 ((500)
16 ((10)
2560 ((60)
4070 ((500)
78900 ((870)
2330 ((38)
a pH ) 4.2, TSS ) 93 mg/L, and TDS ) 332 000 mg/L. Brine was at
room temperature (25 °C) during the pH measurements. b Results provided
by the U.S. DOE National Energy Technology Laboratory.10
it is only reliable at 25 °C; hence the reason for investigation
into customizing the Pitzer database. Although some of the
generic databases that come with the program do account for
higher temperatures, they do not account for the increased ionic
interactions presumed to be observed in such high salinity brines,
and therefore, the use of these generic databases may not give
reliable results.
Characterization of Unreacted Brine and Liquid Brine
Products. Table 2 presents the ICP-AES screening results
conducted with and without acid pretreatment and also includes
the pH, total suspended solids (TSS), and total dissolved solids
(TDS) measurements of the natural untreated OH-1 brine. In
comparison to similar brines used in previous studies, OH-1
(12) Parkhurst, D. L.; Appelo, C. J. User’s guide to PHREEQC (Version
2) - A computer program for speciation, batch-reaction, one-dimensional
transport, and inVerse geochemical calculations, 1999. http://wwwbrr.cr.usgs.gov/projects/GWC_coupled/phreeqc/html/final.html (accessed April
15, 2005).
has a large amount of TDS at 332 000 mg/L.8 Due to the high
measured concentrations of Na ions present within the brine, it
is assumed that the abundance of TDS is due to large amounts
of NaCl dissolved in solution. Also relative to other brines, OH-1
has a significantly low iron concentration, which has previously
been linked to its ability to maintain an elevated adjusted pH.8
The differences between the two data sets reveal that approximately 85% of the metals are in the liquid phase, with
iron as an exception at approximately 55%.
The results of ICP-AES analysis of the sampled and final
liquid products from the high-pressure/high-temperature reactions between CO2 and brine in experiments B through D are
shown in Figure 2. Experiments B through D were all conducted
at a pressure of 600 psi. Also shown in Figure 2 is the evolution
of brine pH throughout the 18-hour reactions. The data show
that the initial pH drop takes place during purging due to the
formation of carbonic acid following the high-pressure-induced
dissolution of CO2 in brine. Due to this rapid drop in pH during
purging, it likely that some reactions occur even before the
system is heated. This may be indicated by the decrease in Ca
in each of the three brines during this time. In regards to the
pH evolution of the system at 150 °C, results do not support a
steady and continued increase in pH following the initial rapid
drop, as was alluded to in our previous study that used a reaction
time of only 6 h.8 These results suggest that their observed
increase was not the evolution of the system toward a higher
pH but was rather a fluctuation in pH before reaching equilibrium.
Figure 2 shows the concentration of various metals as a
function of time for experiments B, C, and D, in which the
initial pH values were 8.97, 7.11, and 4.85, respectively. All
three experiments were conducted at 150 °C and 600 psi, while
176 Energy & Fuels, Vol. 20, No. 1, 2006
Druckenmiller et al.
Figure 3. XRD patterns of the precipitates from experiments A though D. “C” and “H” denote the spectral peaks of calcite and halite, respectively.
experiment A (not included in this figure) was conducted at
150 °C and 1500 psi. The uncertainty in these measurements is
the same for each element as shown in the brine characterization
data in Table 2. The starting concentration of Na for the brine
with the highest initial pH of 8.97 is significantly lower, which
is due to a reduction in solubility for NaCl in brine as a result
of the addition of KOH. In general, the concentration of each
element fluctuates significantly, except for K, which appears
to remain fairly constant, and Fe, which steadily increases. These
large fluctuations also indicate that the system did not reach
steady state. Kaszuba et al. found that their similar experimental
system, which operated at a pH of approximately 5 and at 200
°C and 2900 psi, took 1330 h to reach steady state.13
The overall observed increase in Fe of 170, 110, and 60 ppm
for experiments B, C, and D, respectively, is similar to what
was observed by Kaszuba et al., in which a 4-fold increase in
iron was observed over a 500-h period.13 Additionally, this
increase in iron correlates with that observed by Soong et al.5
Both the initial and continued concentration of K throughout
the reactions is representative of the amount of KOH needed to
adjust the pH of the brines to the desired value.
The beginning pH values of the brine samples also correlate
with the difference between the starting concentrations of Na,
Ca, Mg, and Sr. The brine with the highest starting pH
(experiment B) had the lowest concentration of each of these
elements. Conversely, the brine with the lowest starting pH
(experiment D) had the highest concentration of each element.
This is due to carbonate minerals having a lower solubility at
a higher pH.5
(13) Kaszuba, J. P.; Janecky, D. R.; Snow, M. G. Appl. Geochem. 2003,
18, 1065-1080.
As will be discussed in the following section, experiment B
clearly precipitated calcite, while experiment C showed trace
indications of calcite. In Figure 2, the Ca concentration of each
brine decreases by 3-6% during pressurization and then
increases by 11-17% following heating to 150 °C. The increase
in Ca following heating is opposite of what is expected based
solely on the solubility properties of calcite. As temperature
increases, the solubility of calcite decreases; thus one would
expect the Ca concentration to decrease.14 This same decrease
during pressurization and increase during heating is also the
general trend observed for Na, Mg, and Sr. However, an increase
during heating for Na is expected, because the solubility of NaCl
increases with temperature.15 Finally, the final concentration in
Ca for experiment B, which showed a clear sign of calcite
precipitation, was higher than the initial concentration by 2000
ppm or 10%. This does not correlate with what is expected or
with the results found by Soong et al. in their experiments in
which calcite formation was observed.5 The fact that a decrease
in Ca is not observed in experiment B is likely because the
system has not yet reached a steady state, and thus the decrease
in Ca due to the precipitation of calcite is smaller than the
observed fluctuations in the Ca concentration. The fact that the
Ca concentration during or at the end of an experiment can be
higher than the initial concentration suggests that the excess
Ca comes from that present in the initial unreacted brine in
suspended solids.
Characterization of Solid Products Using XRD. XRD
analysis of the solid products from experiments A through D
(14) Plummer, L. N.; Busenberg, E. Geochim. Cosmochim. Acta 1982,
46, 1011-1040.
(15) Cohen-Adad, R.; Balarew, C.; Tepavitcharova, S.; Rabadjieva, D.
Pure Appl. Chem. 2002, 74, 1811-1821.
Carbon Sequestration Via Calcite Formation in Brine
Energy & Fuels, Vol. 20, No. 1, 2006 177
Figure 5. Low resolution SEM image of the areas of a calcite
precipitate that were used for EDS analysis. Areas labeled 1, 2, and 3
represent the specific areas analyzed using EDS as shown in Figure 6.
Figure 4. SEM images of a calcite precipitate magnified 3000 times.
Images a and b are an SE image and a BSE image, respectively. Note
that the images do not represent the same area in the specimen.
was performed for qualitative characterization. Very small
amounts of precipitates were formed, and therefore, any effort
to quantify produced solids proved difficult. Figure 3 shows
the XRD patterns of the analyzed solids in which the two species
present are halite and calcite, the most tightly packed polymorph
of calcium carbonate. Calcite was clearly identified for experiments A and B, in which the initial pH was approximately 9.0.
These results correlate well with what was observed in our
previous work, in which results showed that calcite was only
formed in experiments with a pH value of approximately 9.0,
as opposed to the lower pH values used in the study.8 Similar
conclusions have been made from previous research by Soong
et al.5 Experiment C with an initial pH of 7.11 reveals a very
small and nearly undetectable presence of calcite. However, the
fact that calcite is present may suggest that, given a longer
reaction time, an increased amount of calcite may precipitate.
Other than calcite, no calcium carbonate polymorphs, such as
aragonite or vaterite, or other carbonates were identified in the
produced solids from these experiments.
Although XRD was not conducted in a manner best suited
for quantitative analysis, a quantitative look at the presence of
calcite can be achieved using the normalized relative intensity
ratio (RIR) method. It should be noted that a refined quantitative
analysis procedure for XRD is achievable with the use of an
internal standard such as quartz. The equation for the normalized
RIR method is
Xa )
Ia
RIRa‚I(rel)a
[∑
]
1
Ix/(RIRx‚I(rel)x)
(1)
where Xa is the weight fraction of phase a, which is the phase
of interest, RIRa is the RIR value for phase a, Ia is the observed
intensity for a peak of phase a, I(rel)a is the relative intensity of
Ia, RIRx is the RIR value for a certain phase, Ix is the observed
intensity for a peak of a certain phase, and I(rel)x is the relative
intensity of Ix.16 RIR values are empirical numbers obtained
for specific phases from the XRD analysis software.
This method can be used, for example, to quantitatively
analyze an XRD profile of a precipitate in which calcite has
been identified by assuming that the entire mass is composed
of only calcite and halite, which is likely very accurate based
on the experimental XRD profiles from experiments A through
C (see Figure 3). Because experiment B resulted in the largest
calcite peak in the XRD analysis, the XRD results from this
experiment are used to calculate the weight percent of calcite
in the precipitate. This allows for an estimate of the amount of
calcite formed during the reaction and thus provides for an
estimate of the total carbon sequestered in mineral form. Once
again, since it is believed that the reaction time of experiment
B was not sufficiently long enough for the system to reach
steady state (i.e., where fluctuations in the solution calcium
concentration are much smaller than the decrease in the calcium
ion concentration resulting from the precipitation of calcite),
estimates of total sequestered carbon will only correspond to
the collected solids and not to the observed calcium concentration trend in the ICP-AES data of Figure 2.
For experiment B, if calcite is denoted as the phase of interest
and halite is the only remaining phase, RIRcalcite is 2.0, Icalcite is
23%, I(rel)calcite is 27%, RIRhalite is 4.40, Ihalite is 63%, and
I(rel)halite is 73%. Solving eq 1 with these values gives an Xcalcite
value of 68%, which is the mass percent of calcite in the
analyzed precipitate. Experiment 18, which began with 95 mL
of brine, precipitated 132 mg of precipitate and therefore
precipitated 90 mg of calcite, which corresponds to 11 mg of
sequestered carbon and 40 mg of avoided CO2. In summary,
an 18-h reaction between CO2 and brine with an initial pH of
approximately 9.0 at 150 °C and 600 psi will sequester 0.42 g
of CO2 per liter of brine. Error in this estimate exists because
small volumes of brine from this specific experiment were
extracted during sampling and, accordingly, it is likely that
precipitates were also extracted. Also, the mesh size of the filter
(16) Wonderling, N. M. Materials Research Institute, The Pennsylvania
State University, University Park, PA, 2002.
178 Energy & Fuels, Vol. 20, No. 1, 2006
Druckenmiller et al.
Figure 6. EDS spectra for the labeled areas of the SEM image shown in Figure 5. The peak labeled Ca′ represents a secondary peak associated
with the primary peak for calcium. The spectrum for area 3 is not presented at the same intensity scale as those for areas 1 and 2.
paper used to separate the precipitates from the brine at the
conclusion of the experiment was not fine enough to maximize
the collection of precipitates. Furthermore, the use of a finer
mesh filter (e.g., 0.5 or 0.2 um) may also improve the analytical
ICP-AES results for the filtrates by removing fine clay or
colloidal materials.
Using a dissolved CO2 density of 40-60 kg CO2/m3 brine,
which is typically used for geologic carbon sequestration
capacity estimates,17 0.42 kg of sequestered CO2/m3 of brine
achieved in an 18-h period would correspond to sequestration
in carbonate form of approximately 0.7-1.1% of the CO2 that
would be dissolved in brine under typical geologic temperature
and pressure conditions. This sequestration estimate may appear
low in the context of a large industrial-scale process developed
to treat a CO2-rich flue gas. However, in these experiments the
brine is not continuously buffered toward a high pH that
promotes carbonate formation. In an industrial process, the
potential to continuously buffer the brine pH clearly would exist,
which would more effectively sequester CO2.
Aside from the aforementioned KOH used in this study as
the primary means to adjust the pH of the brines prior to reaction
with CO2, the potential exists to use a buffer solution to augment
the reaction of CO2 with the brines. In particular, the use of a
0.6 M NaHCO3 + 1 M NaCl “buffer” solution has been studied
in the use of direct carbonation to aid in further sequestration
of CO2.18 While this sodium bicarbonate/salt mixture does not
fit the typical definition of a buffer solution, it is known to
behave similarly to a buffer and is considered to increase the
HCO3- concentration, accelerating the formation of carbonate
ions (see reaction 4 in the Introduction).18,19
Further investigation of the utilization of the NaHCO3/NaCl
mixture with brine is still needed, as preliminary studies suggest
that it cannot be used primarily for pH adjustment. Given that
(17) Bachu, S.; Adams, J. J. Energy ConVers. Manage. 2003, 44, 31513175.
(18) O’Connor, W. K.; Dahlin, D. C.; Rush, G. E.; Dahlin, C. L.; Collins,
W. K. Miner. Metall. Process. 2002, 19, 95-101.
(19) Huijgen, W. J. J.; Comans, R. N. J. Carbon dioxide sequestration
by mineral carbonation: Literature reView; Report ECN SF: ECN-C--03016; Energy Research Centre of The Netherlands: Petten, The Netherlands,
2003.
brine pH has been identified as having the greatest control on
precipitation of calcite, the essentially acidic nature of brine
combined with the weak, basic NaHCO3/NaCl mixture does not
result in favorable pH adjustments. Therefore, a better understanding of the optimal modification of the brine solution
chemistry with the NaHCO3/NaCl mixture is necessary.
SEM/EDS Analysis of Calcite. A solid precipitate from an
experiment reported in our previous work was used for SEM/
EDS imaging and analysis.8 The reaction that produced the
analyzed solid was performed in the same manner and with the
same reactor as experiments A through D discussed in this work.
The reaction was conducted at 1500 psi, 75 °C, and with an
initial pH of approximately 9.0. The precipitate from this specific
experiment was chosen because it was a reaction in which calcite
was clearly present as determined by XRD. Figure 4a shows a
3000×-magnified high resolution secondary electron (SE) image
of the crystals present in the precipitate. The crystals are
irregularly shaped and show no sign of a consistency in size.
One cubic particle is shown as having a width of 12 µm.
Because most of the particles appear smaller than this size, it is
reasonable to assume that a portion of the precipitated solids
was not collected during the brine filtering process, which used
a filter with a particle retention size of 16 µm. The existence of
horizontal bright bands in the image, as well as relatively bright
particle edges, is an effect known as charging. This image defect
is a result of a build of charge in areas under the electron beam
that are of low conductivity. Charging prevents smaller scale
irregularities on the surface of particles from being distinguished.
Figure 4b shows a 3000×-magnified low resolution backscattered electron (BSE) image of the crystals present in the
precipitate. Even though this image was taken at low resolution,
the irregularities on the surface of the larger particles are more
visible because BSE images are less sensitive to charging than
SE images. However, it is visibly evident from the localized
bright regions that specific areas of the specimen did in fact
experience charging, which is likely because these regions had
greater exposure to the electron beam due to their orientation.
In general, in the absence of charging, brighter particles
represent heavier phases and darker particles represent lighter
phases. In the relatively dark regions of the specimen in Figure
Carbon Sequestration Via Calcite Formation in Brine
4b, smaller bright particles are seen on the surface of larger
dark particles. Given that the specimen is not flat, it is difficult
to distinguish regions that are bright due to phase weight or
due to the charging effect as a result of individual particle
orientation or height. However, regions in the specimen that
do not appear oriented in the same manner as the brightest faces
(i.e., oriented in the direction of the beam) still show a textured
surface indicative of smaller particles existing on the surface.
Thus for these relatively dark regions, it is assumed that the
bright small surface particles are calcite, which has a molecular
weight of 100, and the dark larger particles are halite, which
has a molecular weight of 58.4. To enhance this type of analysis
and to eliminate the potential for discrepancy, future work will
aim to minimize the charging effect through SEM improvement
methods such as rapid scanning, which prevents the buildup of
a charge.
Figures 5 and 6 together show the results of the EDS analysis
of an SEM image, which was further magnified to show the
surface of a larger particle similar in scale to those shown in
Figure 4a,b. These figures are meant to illustrate qualitative
differences between the elemental compositions of the three
identified areas of the specimen through a relative comparison
of the height of the Cl peaks with that of the Ca′ peaks, which
are the secondary peaks associated with the calcium primary
peaks. The smaller bright particles are clearly seen in Figure 5.
The spectrum for area 2, which is a bright region, shows that
the Cl peak is small relative to the Ca′ peak. The EDS spectrum
for area 1, which is a dark region, shows that the Cl peak is
approximately the same size as the Ca′ peak. Additionally,
another dark region shown as area 3 has a spectrum in which
both the Na and Cl peaks are much larger than the Ca′ peak.
(In Figure 6, the intensity spectrum for this area is presented at
a smaller scale than for areas 1 and 2 to highlight the clear
difference in height between the Cl peak and the Ca′ peak.)
Therefore, the darker regions are more concentrated in Na and
Cl than the bright regions. These data support the conclusion
that the larger dark particles are halite and the smaller bright
particles are calcite. However, it must once again be noted that
elemental mapping is best performed on a flat specimen, and
therefore, uncertainty exists as to whether these results accurately
address the elemental composition of the areas of interest in
the sample.
Energy & Fuels, Vol. 20, No. 1, 2006 179
Conclusions
Calcite formation was induced at temperatures of 150 °C and
pressures ranging from 600 to 1500 psi by reacting CO2 with a
natural gas well brine of low iron content. The most favorable
initial brine pH for calcite formation was approximately 9.0, as
opposed to the lower values investigated. During the reactions,
the system pH immediately dropped following the introduction
of high-pressure CO2 due to the formation of carbonic acid.
An experimental reaction time of 18 h revealed little change in
system pH following the initial decrease. At the start of the
reactions, the liquid-phase composition of various metals, such
as Na, Ca, Mg, and Sr, was inversely related to the adjusted
pH as minerals have a lower solubility at a high pH. Throughout
the course of the reactions, compositional trends in the metals
most important for mineral carbonate formation were not
identified. Large fluctuations indicate that the system did not
reach steady state. ICP-AES data cannot be used to correlate
changes in brine composition with pH changes and the
precipitation of carbonates using a reaction time of 18 h.
XRD has verified that calcite was most abundantly present
in the solid products for the reactions with an initial pH of
approximately 9.0. No other carbonate minerals, such as
magnesite or dolomite, were detected despite the significant
brine concentrations of magnesium. This finding supports
previous conclusions from our previous work regarding the
importance of pH in inducing calcite formation.8 SEM imaging
suggests that calcite crystals in the precipitate are mostly present
on the surface of larger particles and at sizes much smaller than
12 µm, which is the size of the largest measured particle in the
analyzed specimen.
Feasibility for an industrial-scale operation to sequester carbon
in natural gas well brine is currently limited by the extent that
pH needs to be raised and maintained to promote the rapid
formation of carbonates. As the effects of the various parameters
on carbonate mineral formation are further understood, a refined
assessment of the pH requirements for employed brines will
result.
Acknowledgment. We thank Dr. Yee Soong and Dr. Daniel J.
Fauth at the National Energy Technology Laboratory of the U.S.
DOE for the experimental brine samples and for helpful discussions,
and Dr. Semih Eser at the Energy Institute at The Pennsylvania
State University for the use of the reactor.
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