energies
Article
Comparison of Two Processes Forming CaCO3
Precipitates by Electrolysis
Hyun Sic Park 1 , JunYoung Han 2 , Ju Sung Lee 1 , Kwang-Mo Kim 3 , Hyung Jun Jo 1
and Byoung Ryul Min 1, *
1
2
3
*
Department of Chemical and Biomolecular Engineering, Yonsei University, 262 Seongsanno, Seodaemun-gu,
Seoul 120-749, Korea; [email protected] (H.S.P.); [email protected] (J.S.L.);
[email protected] (H.J.J.)
Department of Chemistry and Chemical Biology, Rensselaer Polytechnic Institute, Troy, NY 12180, USA;
[email protected]
COSMAX #603, Pangyo Inno Valley, 255, Pangyo-ro, Bundang-gu, Seongnam-si, Gyeonggi-do 13486, Korea;
[email protected]
Correspondence: [email protected]; Tel.: +82-2-2123-2757
Academic Editor: Covadonga Pevida García
Received: 28 September 2016; Accepted: 8 December 2016; Published: 13 December 2016
Abstract: As one of the carbon capture and utilization (CCU) technologies, mineral carbonation
which has been introduced to reduce the carbon dioxide (CO2 ) concentration in the atmosphere is
a technology that makes it possible to capture CO2 and recycle byproducts as resources. However,
existing mineral carbonation requires additional energy and costs, as it entails high temperature and
high pressure reaction conditions. This study compared two processes which electrolyze NaCl and
CaCl2 solution to produce CO2 absorbent needed to generate CaCO3 , and which were conducted at
room temperature and pressure unlike existing mineral carbonation. As a result, high-purity calcite
was obtained through Process 1 using NaCl solution, and aragonite and portlandite were obtained in
addition to calcite through Process 2 (two steps) using CaCl2 solution.
Keywords: mineral carbonation; CaCO3 ; electrolysis
1. Introduction
Carbon dioxide (CO2 ) accounts for more than 88% of greenhouse gas emissions and has
contributed more than 55% to the greenhouse effect [1,2]. In order to solve the climate change and
environmental problems caused by CO2 , a number of studies have been carried out actively [1,3,4].
Until now, intensive research has been conducted on carbon capture and storage (CCS)
technologies to store the CO2 underground or on the seabed and carbon capture and utilization (CCU)
technologies to capture CO2 and convert it to be reused. However, it has become known that there
are several problems in the practical application of these techniques [5]. In the case of CCS, excessive
energy consumption throughout the entire process and the possibility of leakage of absorbent have
been pointed out as the biggest problems [6]. CCU technology was introduced as a partial alternative
to these problems. Among the CCU technologies, mineral carbonation is one of the most discussed
techniques for treating large amounts of CO2 . However, the magnesium-based mineral carbonation
process is also complicated and should be conducted in high temperature and high pressure conditions,
so the facility where it is conducted needs to be equipped for the high risk group [7–9]. For this
reason, this study introduces a mineral carbonation process undertaken with less energy by using
electrolysis technology with a porous ceramic membrane as ion channel and separator. In this process,
divalent metal cations are combined with CO2 to form carbonate minerals [10]. Generally, CO2 reacts
with the alkaline absorbent solutions. In the existing carbonation process, magnesium hydroxide was
Energies 2016, 9, 1052; doi:10.3390/en9121052
www.mdpi.com/journal/energies
Energies 2016, 9, 1052
2 of 8
primarily used to produce the alkaline solution, because the Mg-containing minerals contain about
40% Mg2+ , while the Ca-containing minerals contain about 10%–15% Ca2+ [11]. However, the existing
magnesium-based carbonation process has the aforementioned problems. According to Nduagu [12]
and Khoo et al. [13], extracting magnesium oxide or hydroxide from serpentine requires heating to
a temperature of 500 ◦ C or higher, and high temperature reaction conditions are needed again in the
hydration process after extraction [14].
To reduce these problems, our previous study and this study have introduced a process of
electrolyzing chloride salts present in seawater and producing alkaline solution, which is a CO2
absorbent. Our process does not extract oxides and hydroxides from rocks, but uses electrolysis
at room pressure and room temperature to obtain hydroxides. More specifically, in our previous
report [15], a calcium carbonation process using NaCl solution was introduced. The entire process
is divided into three steps: electrolysis of NaCl (sodium hydroxide formation), reaction with CO2
(sodium bicarbonate formation), and carbonation (generation of calcium carbonate). Through this
process, we were able to obtain CaCO3 in the form of calcite in a stable manner. However, since NaCl
solution is used for NaCl process, Ca2+ should be added to the carbonation process in order to replace
Na+ . In addition, the process of precipitating the carbonate after this substitution process is complex
and requires considerable time.
Therefore, in this study, in order to simplify process and increase energy efficiency, CaCl2 in
seawater was used to make alkaline solution, so that the process of Ca2+ addition was eliminated and
the process was reduced to two steps, unlike in our previous study. Since the substitution process is
not required, the carbonation process is simpler than the NaCl solution process. Additionally, it was
expected that more products could be obtained because of the high content of calcium ions contained
in the carbonation process. Precipitates would also be formed more quickly through this process, and
the consumption of energy would be decreased. This present work focuses on the comparison of the
characterization and yield of final products for our previous and proposed processes.
2. Materials and Methods
2.1. Materials
NaCl (99.5%) and CaCl2 (95.9%) were purchased from Samchun Chemical Korea and Kanto Co.,
Tokyo, Japan, respectively, and used as received. Pure CO2 (99.0%) gas was purchased from Samheung,
Gyeonggi-do, Korea, and was used without any pretreatment. Pure CaCO3 (≥99.0%, Sigma-Aldrich,
St. Louis, MO, USA) was used to compare with the final CaCO3 product from the two processes.
Ultra-pure water prepared from AquaMax 311 (YL Instrument Co. Ltd, Anyang, Korea) was used in
all processes.
2.2. Electrolysis Device
The electrolysis device is shown in Figure 1. Each of the units is the same as those used before [15].
The anode (carbon) is located in the center of a round ceramic membrane. The electrodes of the cathode
(stainless) were installed at the same distance around the anode. The mean pore size of the porous
ceramic membrane was 0.2 µm, which acts as an ion channel when the electrolytic process is started.
In other words, ion separation occurs between the inner and outer reactor during the electrolysis
process. After the electrolytic process, fine CO2 bubbles (100–400 µm) were injected from the generator
(ANGEL AQUA, Changwon, Korea). At this time, this ceramic wall plays as an ion separator.
Energies 2016, 9, 1052
3 of 8
Energies 2016, 9, 1052
3 of 8
Figure 1. Schematic diagram of the electrolysis device.
Figure 1. Schematic diagram of the electrolysis device.
2.3. NaCl Solution Process: Process 1
2.3. NaCl Solution
Process:
In the
case ofProcess
the NaCl1 solution
process (Process 1), the experiment to generate CaCO3 was
carried out in three steps, preparing 2%–6% NaCl solution based on the concentration of coastal
In the case
of the
solutionDepending
process (Process
the experiment
to generate
CaCO
seawater
andNaCl
deep seawater.
on 1–4 A of1),
current
intensity, electrolysis
of NaCl
solution
3 was carried
required
to react for
10–15 min
at 25solution
°C and 1 bar,
and voltage
about 8–24 V. Characterization
out in threewas
steps,
preparing
2%–6%
NaCl
based
on thewas
concentration
of coastal seawater and
of the product was reported in our earlier paper [15], and chemical reaction of each step was
deep seawater.
Depending on 1–4 A of current intensity, electrolysis of NaCl solution was required to
summarized as follows:
◦
react for 10–15 min at 25 C and 1 bar, and voltage was about 8–24 V. Characterization of the product
Step 1:
was reported in our earlier paper [15], and chemical reaction of each step was summarized as follows:
electrolysis
Step 1:
2NaCl (aq) + 2H2 O →
Step 2:
2Na+ + 2OH− + H2 (g)↑ + Cl2 (g)↑
(1)
electrolysis
+
−
2NaCl (aq) + 2H2 O −−−−−−
+ 2OH
+-H2 (g) ↑ +Cl2 (g) ↑
-→ 2Na
+
+
Na + OH + CO2 (g) ⟶ Na + HCO3
Step 2:
(2a)
(1)
pH control
+ −
+
−
12
2+3 OH
+(aq)
CO→2from
Na
HCO
(g6.5) to→
Na+Na
+ HCO
+ NaOH
2Na+ ++
CO
3 + H32 O (l)
Step 3:
(2b)
pH control from 6.5 to 12
Na+ + HCO3− + NaOH
(2-aq) −−−−−−−−−−−−−→ 2Na+ + CO23− + H2 O (l)
+
2Na + CO3 + CaCl2 (aq)⟶ CaCO3 (s)↓ + 2NaCl (aq)
(3)
(2a)
(2b)
Step 3:
+
2− 2
2.4. CaCl2 Solution
Process
2NaProcess:
+ CO
(3)
3 + CaCl2 (aq) → CaCO3 (s) ↓ +2NaCl (aq)
In the case of the CaCl2 solution process (Process 2), the experiment to generate CaCO3 was
performed
in two steps,
as 0.5%–2%
CaCl2 is based on the concentration of seawater. All other
2.4. CaCl2 Solution
Process:
Process
2
parameters are the same as in Process 1. The reaction time was stopped at the point where there is
In the almost
case of
the CaCl
solution
(Process
2), the experiment
to of
generate
CaCO3 was
no variation
in 2pH,
at aroundprocess
11.6. Under
the same conditions,
the reaction time
electrolysis
2 took 5 as
min0.5%–2%
less than in CaCl
Process2 1.is based on the concentration of seawater. All other
performed ininProcess
two steps,
Step 1, an alkaline calcium hydroxide solution used as an absorbent of CO2 was produced
parameters areInthe
same as in Process 1. The reaction time was stopped at the point where there is
through electrolysis, and the reaction is as follows:
almost no variation in pH, at around 11.6. electrolysis
Under the same conditions, the reaction time of electrolysis
(4)
CaCl2than
(aq) +in2H
Ca2+ + 2OH- + H2 (g)↑ + 2Cl2 (g)↑
2 O (l) → 1.
in Process 2 took 5 min less
Process
2, calcium
hydroxide
solution was
formedused
by a ceramic
membrane atofthe
negative
In Step 1, In
anStep
alkaline
calcium
hydroxide
solution
as an absorbent
CO
2 was produced
electrode side after electrolysis. CaCO3 was precipitated by blowing CO2 at constant flow rate (2 L·min−1).
through electrolysis, and the reaction is as follows:
The precipitation reaction is as follows:
pH control
electrolysis
2+
from 6.5 to 12
CaCl
(aq-3)++OH
2H
O (l) (aq)
−−−−−−
→ Ca
2OH− + H+ (NaOH
g) ↑ +
2Cl2 (g) ↑
Ca2+ +2HCO
+ 2NaOH
→
CaCO+
(aq)
3 (s)↓ + H2 O (l) 2
(5a)
(4)
In Step 2, calcium hydroxide solution was formed by a ceramic membrane at the negative electrode
side after electrolysis. CaCO3 was precipitated by blowing CO2 at constant flow rate (2 L·min−1 ).
The precipitation reaction is as follows:
pH control from 6.5 to 12
Ca2+ + HCO3− + OH− + NaOH (aq) −−−−−−−−−−−−−→ CaCO3 (s) ↓ +H2 O (l) + NaOH (aq) (5a)
pH 12 ↑
Ca2+ + OH− + NaOH (aq) −−−−→ Ca(OH)2 (s) ↓ +Na+
A side reaction as in Equation (5b) occurs when the pH value is higher than 12.
(5b)
Energies 2016, 9, 1052
4 of 8
Energies 2016, 9, 1052
pH 12↑
Ca2+ + OH- + NaOH (aq) →
2.5. Characteristics of Precipitates
Ca(OH)2 (s)↓ + Na+
4 of 8
(5b)
A side reaction as in Equation (5b) occurs when the pH value is higher than 12.
The final product CaCO3 was obtained by GF/C film (Whatman® Glass microfiber filters,
2.5. Clifton,
Characteristics
of Precipitates
Whatman,
NJ, USA)
filtration and drying for 24 h at 80 ◦ C. Fourier transform infrared (FT-IR)
® Glass microfiber filters,
final product
CaCO3 was100
obtained
by GF/C film (Whatman
spectra wereThe
obtained
with a Spectrum
FT-IR Spectrometer,
PerkinElmer
connected attenuated total
Whatman,
Clifton,
NJ,
USA)
filtration
and
drying
for
24
h
at
80
°C.
Fourier
transform
(FT- using a
reflection (ATR) tool (PerkinElmer, Norwalk, CT, USA). X-ray diffraction (XRD)
wasinfrared
measured
IR) spectra were obtained with a Spectrum 100 FT-IR Spectrometer, PerkinElmer connected
Rigaku Ultima IV (XRD, Ultima IV, RIGAKU, Tokyo, Japan) using a Cu Kα X-ray 40 kV/30 mA, and
attenuated total reflection
(ATR) tool (PerkinElmer, Norwalk, CT, USA). X-ray diffraction (XRD) was
the range
was 2θ = 15◦ –70◦ . The shape and particle size of the fully dried samples were investigated
measured using a Rigaku Ultima IV (XRD, Ultima IV, RIGAKU, Tokyo, Japan) using a Cu Kα X-ray
by field40
emission
scanning
electron
(FE-SEM,
JEOL-7800F,
Japan).
kV/30 mA,
and the range
was microscopy
2θ = 15°–70°. The
shape and
particle sizeTokyo,
of the fully
dried samples
were investigated by field emission scanning electron microscopy (FE-SEM, JEOL-7800F, Tokyo,
3. Results
Japan).
3.1. Fourier
Transform Infrared Analysis
3. Results
Figure
2 shows
the IRInfrared
spectra
measured in the 1800–600 cm−1 range of the respective samples.
3.1. Fourier
Transform
Analysis
All CaCO3 products from Process 1 and Process 2 were confirmed
as calcite, and the peaks were 1418,
Figure
the IR spectra measured in the 1800–600 cm−1 range of the respective samples.
−12. shows
876, andAll713
cm
Interestingly,
the
IR
spectrum
of
CaCO
obtained
Process
2 was
observed
3
CaCO3 products from Process 1 and Process 2 were confirmed
as calcite,from
and the
peaks were
1418,
−
1
−1
−1. Interestingly,
for new876,
peaks.
Thecm
value
of the peak
1500–1400
region2 and
the 700–650
and 713
the observed
IR spectruminofthe
CaCO
3 obtained cm
from Process
was observed
for cm
−
1
−
−1 region
−1 region
value of3 the
peak observed
in the 1500–1400
and the
700–650 cm
region isnew
thepeaks.
peakThe
of CaCO
as aragonite
(respectively,
1465cmcm
, C–O,
stretching
mode;
692 cm 1 ,
−1
−1
−1 region
is the peakbending),
of CaCO3 asand
aragonite
(respectively,
1465 cm cm
, C–O,
stretching
692 cm
, C–O,
out- only as
C–O, out-plane
the peak
in the 850–800
is mode;
the value
of the
peak
plane bending), and the peak
in the 850–800 cm−1 region is the value of the peak only as vaterite and −1
−
1
vaterite and aragonite (848 cm , C–O, in-plane bending) [16,17]. In addition, the peak (1080 cm ) in
aragonite (848
cm−1, C–O, in-plane bending) [16,17]. In addition, the peak (1080 cm−1) in the 1100–1000
the 1100–1000
cm−1 region is the stretching mode of the O–C–O functional group.
−1
cm region is the stretching mode of the O–C–O functional group.
Figure 2. Fourier transform infrared (FT-IR) spectrum of the precipitated particles through Process 1
Figure 2. Fourier transform infrared (FT-IR) spectrum of the precipitated particles through
Process 1
and Process 2. (a) N-2; (b) N-3; (c) N-4; (d) N-5; (e) N-6; and (f) C-0.5 (range of 1800–600 cm−1). Product
−1 ). Product
and Process
2.
(a)
N-2;
(b)
N-3;
(c)
N-4;
(d)
N-5;
(e)
N-6;
and
(f)
C-0.5
(range
of
1800–600
cm
particles formed using CaCl2 0.5% solution are termed as C-0.5, while product particles formed using
particlesNaCl
formed
usingas
CaCl
0.5% solution
are concentration.
termed as C-0.5, while product particles formed using
are termed
N-X,2 where
X is the NaCl
NaCl are termed as N-X, where X is the NaCl concentration.
3.2. X-ray Diffraction Analysis
XRD patterns
were analyzed for more accurate compositional analysis of the generated particles
3.2. X-ray Diffraction
Analysis
and the results are shown in Figure 3. The product of Process 1 is that characteristic peaks of calcite
XRD
patterns
were analyzed
for more
accurate
analysis
of theofgenerated
particles
only
appear through
our previous
studies.
In XRD compositional
pattern analysis of
the products
Process 2, it
and thewas
results
are
shown
in
Figure
3.
The
product
of
Process
1
is
that
characteristic
peaks
of
possible to confirm the two-theta values for aragonite (26.3, 27.2) and for Ca(OH)2 (34.1, 47.1), as calcite
well
the two-theta
value for calcite
(29.3,
[18]. analysis of the products of Process 2, it was
only appearas
through
our previous
studies.
In 39.3,
XRD43.1)
pattern
According
to
the
research
of
Ma
et
al.
[19],
crystalline
forms
CaCO
3 are greatly influenced by
possible to confirm the two-theta values for aragonite (26.3,
27.2)ofand
for
Ca(OH)2 (34.1, 47.1), as well
the pH of the water soluble matrix (WSM), and it has been reported that aragonite crystal forms can
as the two-theta value for calcite (29.3, 39.3, 43.1) [18].
According to the research of Ma et al. [19], crystalline forms of CaCO3 are greatly influenced
by the pH of the water soluble matrix (WSM), and it has been reported that aragonite crystal forms
can be obtained at low pH levels (pH 5.5–9.0). For Process 2, after forming a Ca(HCO3 )2 , CaCO3
was precipitated by adjusting the pH in the last step. Thus, as reported in the case where the
precipitation starts from a low pH level, aragonite crystals can be obtained. However, as the pH
level gradually increases, it is determined that crystalline calcite was also generated, as well as
Energies 2016, 9, 1052
5 of 8
Energies
2016, 9, 1052
5 of 8
be2016,
obtained
Energies
9, 1052 at low pH levels (pH 5.5–9.0). For Process 2, after forming a Ca(HCO3)2, CaCO3 was 5 of 8
precipitated by adjusting the pH in the last step. Thus, as reported in the case where the precipitation
be
obtained
lowpH
pHlevel,
levelsaragonite
(pH 5.5–9.0).
Forcan
Process
2, afterHowever,
forming aasCa(HCO
2, CaCO
3 was
starts
from aatlow
crystals
be obtained.
the pH3)level
gradually
precipitated
by
adjusting
the that
pH incrystalline
the −
last step.
Thus,
as reported
in the case
where
thearagonite
precipitation
increases,
it
is
determined
calcite
was
also
generated,
as
well
as
[20]. into
aragonite [20]. Furthermore, the Ca(OH) which exists in the form of hydroxide ions is converted
− which
starts
from a low
pH level,
aragonite
crystals
can be
obtained.
However,
as the pH
level
gradually
Furthermore,
the
Ca(OH)
exists
in
the
form
of
hydroxide
ions
is
converted
into
Ca(OH)
2 and
Ca(OH)2 and precipitates by pH control [21]. For Process 1, after adjusting the pH to 12 or higher by
increases,
it by
is pH
determined
that For
crystalline
calcite
was also the
generated,
as by
aragonite
[20].
precipitates
control [21].
Process 1,
after adjusting
pH to 12asorwell
higher
using NaOH
usingFurthermore,
NaOH solution,
precipitation
wasinstarted
through substitution
reaction with
the calcium
ions.
−
the Ca(OH)
exists
thesubstitution
form of hydroxide
converted
Ca(OH)
solution, precipitation
waswhich
started
through
reactionions
withis the
calciuminto
ions.
Thus, 2ifand
the
Thus,precipitates
if the precipitation
is
started
at
a
high
pH,
it
is
possible
to
obtain
only
the
hexahedral
form
of
byispH
control
ForpH,
Process
after adjusting
theonly
pH to
or higher byform
using
precipitation
started
at [21].
a high
it is 1,possible
to obtain
the12hexahedral
ofNaOH
calcite
calcite
precipitation
[20,22].
solution,
precipitation
precipitation
[20,22]. was started through substitution reaction with the calcium ions. Thus, if the
precipitation is started at a high pH, it is possible to obtain only the hexahedral form of calcite
precipitation [20,22].
Figure
3. The
X-ray
diffraction(XRD)
(XRD)spectra
spectra of
particles
from
Process
2. 2.
Figure
3. The
X-ray
diffraction
of the
theproduced
produced
particles
from
Process
Figure 3.Scanning
The X-rayElectron
diffraction
(XRD) spectra
the produced particles from Process 2.
3.3. Field Emission
Microscopy
ImageofAnalysis
3.3. Field Emission Scanning Electron Microscopy Image Analysis
FE-SEM
wasScanning
used to confirm
other
characteristics
of other formed particles. Figure 4 is the SEM
3.3.
Field
Emission
Electron
Microscopy
Image Analysis
FE-SEM
used
to confirm
other
of other
formed
particles.
Figure
4 is the SEM
image of was
formed
particles
pictured
at characteristics
×10,000 magnification.
From
Figure
4, numerous
differences
was
topictured
confirm
other
characteristics
of other
formed
particles.
4 is thedifferences
SEM
imagewere
of FE-SEM
formed
particles
at ×
From
Figure
4,Figure
numerous
observed
in used
the size
and shape
of10,000
formedmagnification.
particles.
The size
of formed
particles
obtained
from
of1formed
particles
at formed
×10,000
magnification.
From
Figure
4, numerous
differences
wereimage
observed
the
size
andpictured
shape
of
particles.
The
size of
formed
particles
obtained
Process
isin
3–4
μm,
regardless
of the
concentration
of feed
solution,
and
the
particles
formed
werefrom
were
observed
in
the
size
and
shape
of
formed
particles.
The
size
of
formed
particles
obtained
from
cubic
calcite.
On
the
other
hand,
formed
particles
obtained
from
Process
2
were
smaller
than
1
μm,were
Process 1 is 3–4 µm, regardless of the concentration of feed solution, and the particles formed
Process
1
is
3–4
μm,
regardless
of
the
concentration
of
feed
solution,
and
the
particles
formed
were
it was
observed
thathand,
several
shapesparticles
of particles
exist. When
CaCO32crystals
are formed
cubicand
calcite.
On
the other
formed
obtained
fromthe
Process
were smaller
thanat1 µm,
cubic pH,
calcite.
On the
other hand,
formed particles
obtained
from size
Process
smaller than 1 μm,
crystal
nucleation
proceeds
and affects
[19].2 were
and ithigh
was observed
that
several
shapesrapidly
of particles
exist. particle
When the
CaCO
crystals
are
formed
at high
3
and it was observed that several shapes of particles exist. When the CaCO3 crystals are formed at
pH, crystal nucleation proceeds rapidly and affects particle size [19].
high pH, crystal nucleation proceeds rapidly and affects particle size [19].
(a)
(b)
(c)
(a)
(b)
(c)
(d)
(e)
(f)
(d)
(e)
(f)
Figure 4. Scanning electron microscopy (SEM) image of the precipitated particles through Process 1
and Process 2 (×10,000 magnification). (a) N-2; (b) N-3; (c) N-4; (d) N-5; (e) N-6; and (f) C-0.5.
Energies 2016, 9, 1052
6 of 8
Figure 4. Scanning electron microscopy (SEM) image of the precipitated particles through Process 1 6 of 8
and Process 2 (×10,000 magnification). (a) N-2; (b) N-3; (c) N-4; (d) N-5; (e) N-6; and (f) C-0.5.
Energies 2016, 9, 1052
According to
of Westin
and and
Rasmuson
[23], the
growth
rate of calcite
was
4.29 ×was
10−6
According
tothe
theresearch
research
of Westin
Rasmuson
[23],
the growth
rate of
calcite
−
6
−
2
−
1
−
6
−
2
−
1
−2
−1
−6
−2
−1
mol·×
m 10
·s and
wasof1.31
× 10 mol·
·s ×
under
conditions.
Inreaction
the case
·s under
4.29
molthat
·m of·saragonite
and that
aragonite
wasm1.31
10 the
molsame
·m reaction
the same
of
Process
2,
the
intensity
of
calcite
peak
was
strong,
as
shown
in
IR
and
XRD
spectra.
It
was
assumed
conditions. In the case of Process 2, the intensity of calcite peak was strong, as shown in IR and
that numerous
whose that
growth
rates were
faster
in three
timesrates
werewere
formed.
Additionally,
in
XRD
spectra. It calcites
was assumed
numerous
calcites
whose
growth
faster
in three times
the case
of Process
2, in which
crystalline
product2,isinformed
an additional
product
through
were
formed.
Additionally,
in athe
case of Process
which with
a crystalline
product
is formed
witha
side
reaction, product
it is confirmed
thereaction,
size of the
is that
smaller
compared
to the products
of
an
additional
throughthat
a side
it is particles
confirmed
the size
of the particles
is smaller
Process 1, to
due
toproducts
a hindrance
of growth
between
particles.of growth between particles.
compared
the
of Process
1, due
to a hindrance
3.4.
3.4. Conversion
Conversion Yield
Yield of
of Precipitate
Precipitate
Figure
Figure 55 shows
shows the
the average
average conversion
conversion rate
rate of
of formed
formed precipitates
precipitates in
in two
two processes.
processes. The
The rates
rates
for
the
two
processes
were
calculated
by
the
rate
of
theoretical
value
and
measured
value
for
overall
for the two processes were calculated by the rate of theoretical value and measured value for overall
2+ contained
Ca
contained in
in used
used solution
solution in
in the final step, which is the precipitation procedure. In Process 1,
Ca2+
conversion
conversion rate
rate increased
increased as
as the
the concentration
concentration of
of feed
feed solution
solution increased.
increased. However,
However, itit was
was confirmed
confirmed
that
the
average
conversion
rate
was
77.7%
in
5%
solution,
and
it
was
decreased
in
6%
solution.
that
conversion rate was 77.7% in 5% solution, and it was decreased in 6% solution.
In
In
Process
the
averageconversion
conversionrate
rateofofprecipitates
precipitateswas
was90.0%
90.0%in
in0.5%
0.5% solution,
solution, which is more than
Process
2, 2,
the
average
than
10%
10% higher
higher than
than the
the highest
highest value
value shown
shown in
inProcess
Process1.1. The
The reason
reasonfor
forthis
thisdifference
differenceisisthat
thatNa
Na++ and
and
2+
2
−
2+
2−
Ca
Ca of the formed absorbents reacted
reacted with
with different
differentamount
amountofofCO
CO3 3 ,, although the overall amount
amount
of
CO22 gas
gaswas
wasthe
thesame
sameininallall
experiments.
In addition,
reason
conversion
of supplied CO
experiments.
In addition,
thethe
reason
thatthat
conversion
rate rate
was
was
decreased
in
high
concentration
CaCl
solution
was
because
precipitates
were
not
dramatically
decreased in high concentration CaCl2 solution
was because precipitates were
dramatically
2
2+.
increased
increased compared
compared to
to the
the increased
increasedamount
amountof
ofCa
Ca2+
Figure 5.
5. Conversion
Conversion rate
rate of
of precipitates
precipitates for
for each
each concentration
concentration feed
feed solution
solution used in
in Process
Process 11 and
and
Figure
Process
2.
Process 2.
4. Conclusions
Conclusions
4.
In this
this study,
study,attempts
attemptsto
to minimize
minimize the
the problems
problems of
of existing
existing CCS
CCS techniques
techniques were
were made
made by
by using
using
In
+ and Ca2+
+
2+
Na
contained
in
sea
water,
and
two
processes
for
capturing
CO
2
and
producing
carbonate
Na and Ca contained in sea water, and two processes for capturing CO2 and producing carbonate
werecompared
comparedand
and
analyzed.
process
themethods
two methods
have presented
can be
were
analyzed.
TheThe
process
of theoftwo
that wethat
havewe
presented
can be compared
compared
in three
in
three major
ways.major ways.
First,
it
is
possible
be approached
from
the perspective.
energy perspective.
Bothform
processes
form
First, it is possible
to betoapproached
from the
energy
Both processes
precipitates
precipitates
in
a
thermodynamically
stable
form.
Additionally,
the
overall
process
uses
reaction
in a thermodynamically stable form. Additionally, the overall process uses reaction conditions of
conditions
of room
temperature
anddemonstrates
pressure. This
CaCO3quickly
can be and
produced
room
temperature
and
pressure. This
thatdemonstrates
CaCO3 can bethat
produced
stably
quickly
and
stably
using
less
energy
than
conventional
methods
requiring
high-temperature
and
using less energy than conventional methods requiring high-temperature and high-pressure reaction
high-pressure
reaction
whichstructural
was proven
through
structural
analysis using
FT-IR
conditions,
which
was conditions,
proven through
analysis
using
FT-IR spectra.
Further,
the spectra.
energy
Further,
the
energy
consumption
of
Process
2
was
small
compared
to
Process
1,
through
consumption of Process 2 was small compared to Process 1, through simplification to a two-step
simplification
to a two-step
reduced process
of electrolysis
reaction. Inprocess
addition,
process
and reduced
processprocess
time ofand
electrolysis
reaction.time
In addition,
the carbonation
is the
an
carbonationreaction,
process and
is andoes
exothermic
reaction,
and does
not [24].
require additional energy [24].
exothermic
not require
additional
energy
Energies 2016, 9, 1052
7 of 8
Second, it is possible to be approached by added value of formed precipitates. The result of
SEM imaging and XRD pattern analysis of the precipitate produced showed that CaCO3 produced by
Process 1 could be used as chemical material without any purification process after it was produced
with high purity calcite. This has also been proven through our previous studies. On the other hand,
in the precipitate produced by Process 2, aragonite—another crystal form of CaCO3 —was identified
along with portlandite. Aragonite precipitate is a raw material used widely, not only in plastic, rubber,
paint, and paper, but also in feeds and pharmaceuticals. If only aragonite CaCO3 is to be formed,
50–70 ◦ C temperature should be used as reaction conditions [25]. It is also possible to weakly control
the size of precipitates by controlling pH. Through this process, purity and shape control of aragonite
can be applied as raw material in various fields.
Finally, it is possible to be approached from the perspective of CO2 reduction, as it relates to global
warming. In this perspective, it was assumed that Process 2, for which the reaction conversion rate was
10% higher than Process 1, captured more CO2 . Furthermore, the final product of Process 2 can process
more CO2 than that of Process 1, and can be used as raw material for construction materials [26].
Thus, using Process 2 makes it possible to obtain a precipitate with a high added value more
quickly and easily than Process 1. In addition, this study was processed as batch reaction. If precipitates
are produced by continuous reactor, solutions left after reaction can be used again in the reaction by
reprocessing. Therefore, it is possible to minimize environmental problems caused by waste liquor
formed in the process. In addition, natural CaCO3 formation generally requires long reaction time,
as stated in the previous study [15]. On the contrary, both processes we have proposed can reduce the
reaction time to form CaCO3 .
As mentioned earlier, conventional mineral carbonation is subject to the process of extracting
the reaction components. According to Kakizawa et al. [27], it consumes 1.27 MW of energy in the
process for extracting calcium ion from calcium silicate. However, in our studies, extraction process
was excluded. In actual seawater application, a membrane enrichment system would be used, which
would feed the reaction components through separation and concentration by the membrane.
More research is needed, but we hope to contribute to the reduction of atmospheric CO2 through
carbonation of CO2 based on these considerations.
Acknowledgments: This work was supported by the Human Resources Program in Energy Technology of the
Korea Institute of Energy Technology Evaluation and Planning (KETEP), granted financial resource from the
Ministry of Trade, Industry & Energy, Republic of Korea. (No. 20154010200810).
Author Contributions: Hyun Sic Park, JunYoung Han and Byoung Ryul Min conceived and designed the
experiments; Hyun Sic Park performed the experiments; Hyun Sic Park and JunYoung Han analyzed the data;
Ju Sung Lee, Kwang-Mo Kim and Hyung Jun Jo contributed reagents/materials/analysis tools; Hyun Sic Park
wrote first draft; Byoung Ryul Min revised the paper.
Conflicts of Interest: The authors declare no conflict of interest.
References
1.
2.
3.
4.
5.
6.
Yu, K.M.K.; Curcic, I.; Gabriel, J.; Tsang, S.C.E. Recent advances in CO2 capture and utilization. ChemSusChem
2008, 1, 893–899. [CrossRef] [PubMed]
West, T.O.; Pena, N. Determining thresholds for mandatory reporting of greenhouse gas emissions.
Environ. Sci. Technol. 2003, 37, 1057–1060. [CrossRef] [PubMed]
Baciocchi, R.; Corti, A.; Costa, G.; Lombardi, L.; Zingaretti, D. Storage of carbon dioxide captured in
a pilot-scale biogas upgrading plant by accelerated carbonation of industrial residues. Energy Procedia 2011,
4, 4985–4992. [CrossRef]
Holloway, S. Underground sequestration of carbon dioxide a viable greenhouse gas mitigation option.
Energy 2005, 30, 2318–2333. [CrossRef]
Siefert, N.S.; Litster, S. Exergy and economic analyses of advanced IGCC-CCS and IGFC-CCS power plants.
Appl. Energy 2013, 107, 315–328. [CrossRef]
Damen, K.; Faaij, A.; Turkenburg, W. Health, safety and environmental risks of underground CO2
storage–overview of mechanisms and current knowledge. Clim. Chang. 2006, 74, 289–318. [CrossRef]
Energies 2016, 9, 1052
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
8 of 8
Lackner, K.S.; Butt, D.P.; Wendt, C.H. Progress on binding CO2 in mineral substrates. Energy Convers. Manag.
1997, 38, S259–S264. [CrossRef]
Fagerlund, J.; Nduagu, E.; Zevenhoven, R. Recent developments in the carbonation of serpentinite derived
Mg(OH)2 using a pressurized fluidized bed. Energy Procedia 2011, 4, 4993–5000. [CrossRef]
Fagerlund, J.; Nduagu, E.; Romao, I.; Zevenhoven, R. CO2 fixation using magnesium silicate minerals part 1:
Process description and performance. Energy 2012, 41, 184–191. [CrossRef]
Oelkers, E.H.; Gislason, S.R.; Matter, J. Mineral carbonation of CO2 . Elements 2008, 4, 333–337. [CrossRef]
Goff, F.; Lackner, K. Carbon dioxide sequestering using ultramafic rocks. Environ. Geosci. 1998, 5, 89–101.
[CrossRef]
Nduagu, E. Mineral Carbonation: Preparation of Magnesium Hydroxide [Mg(OH)2 ] from Serpentinite Rock.
Master’s (Eng) Thesis, Åbo Akademi University, Åbo, Finland, 2008.
Khoo, H.; Bu, J.; Wong, R.; Kuan, S.; Sharratt, P. Carbon capture and utilization: Preliminary life cycle CO2 ,
energy, and cost results of potential mineral carbonation. Energy Procedia 2011, 4, 2494–2501. [CrossRef]
Stasiulaitiene, I.; Fagerlund, J.; Nduagu, E.; Denafas, G.; Zevenhoven, R. Carbonation of serpentinite rock
from Lithuania and Finland. Energy Procedia 2011, 4, 2963–2970. [CrossRef]
Park, H.S.; Lee, J.S.; Han, J.; Park, S.; Park, J.; Min, B.R. CO2 fixation by membrane separated NaCl electrolysis.
Energies 2015, 8, 8704–8715. [CrossRef]
Tlili, M.; Amor, M.B.; Gabrielli, C.; Joiret, S.; Maurin, G.; Rousseau, P. Characterization of caco3 hydrates by
micro-raman spectroscopy. J. Raman Spectrosc. 2002, 33, 10–16. [CrossRef]
Wu, G.; Wang, Y.; Zhu, S.; Wang, J. Preparation of ultrafine calcium carbonate particles with micropore
dispersion method. Powder Technol. 2007, 172, 82–88. [CrossRef]
McMurdie, H.F.; Morris, M.C.; Evans, E.H.; Paretzkin, B.; Wong-Ng, W.; Ettlinger, L.; Hubbard, C.R. Standard
X-ray diffraction powder patterns from the jcpds research associateship. Powder Diffr. 1986, 1, 64–77.
[CrossRef]
Ma, Y.; Gao, Y.; Feng, Q. Effects of pH and temperature on CaCO3 crystallization in aqueous solution with
water soluble matrix of pearls. J. Cryst. Growth 2010, 312, 3165–3170. [CrossRef]
Stocks-Fischer, S.; Galinat, J.K.; Bang, S.S. Microbiological precipitation of CaCO3 . Soil Biol. Biochem. 1999, 31,
1563–1571. [CrossRef]
Ren, L.; Zhang, Y.; Bian, Y.; Liu, X.; Liu, C. Investigation of quartz flotation from decarburized
vanadium-bearing coal. Physicochem. Probl. Miner. Process. 2015, 51, 755–767.
Han, Y.S.; Hadiko, G.; Fuji, M.; Takahashi, M. Crystallization and transformation of vaterite at controlled pH.
J. Cryst. Growth 2006, 289, 269–274.
Westin, K.-J.; Rasmuson, A.C. Crystal growth of aragonite and calcite in presence of citric acid, DTPA, EDTA
and pyromellitic acid. J. Colloid Interface Sci. 2005, 282, 359–369. [CrossRef] [PubMed]
Wang, X.; Maroto-Valer, M.M. Integration of CO2 capture and mineral carbonation by using recyclable
ammonium salts. ChemSusChem 2011, 4, 1291–1300. [CrossRef] [PubMed]
Colfen, H.; Qi, L. A systematic examination of the morphogenesis of calcium carbonate in the presence of
a double-hydrophilic block copolymer. Chemistry 2001, 7, 106–116. [CrossRef]
Iizuka, A.; Fujii, M.; Yamasaki, A.; Yanagisawa, Y. Development of a new CO2 sequestration process utilizing
the carbonation of waste cement. Ind. Eng. Chem. Res. 2004, 43, 7880–7887. [CrossRef]
Kakizawa, M.; Yamasaki, A.; Yanagisawa, Y. A new CO2 disposal process via artificial weathering of calcium
silicate accelerated by acetic acid. Energy 2001, 26, 341–354. [CrossRef]
© 2016 by the authors; licensee MDPI, Basel, Switzerland. This article is an open access
article distributed under the terms and conditions of the Creative Commons Attribution
(CC-BY) license (http://creativecommons.org/licenses/by/4.0/).