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Steel et al 2013_Conversion of CO2 into mineral carbonates
Conversion of CO2 into mineral carbonates using a regenerable buffer
to control solution pH
Karen M. Steel ⇑
, Kimia Alizadehhesari, Reydick D. Balucan, Bruno Bašic´
School of Chemical Engineering, University of Queensland, St. Lucia, Queensland 4072, Australia
h i g h l i g h t s
Use of a regenerable buffer (tertiary amine) is studied to enable mineral carbonation.
Tertiary amines complex protons to give a pH of 8.2 which enables MgCO3 precipitation.
Increasing temperature from 18 to 85 °C decreases pH by 2.5 pH units.
Higher temperatures 85 °C might enable low pHs needed for Mg–silicate dissolution.
a r t i c l e i n f o
Received 25 October 2012
Received in revised form 14 March 2013
Accepted 16 April 2013
Available online 30 April 2013
a b s t r a c t
The barrier that is currently stalling the rapid conversion of magnesium silicate deposits into magnesium
carbonate as method for storing CO2 is considered to be the difference in pH needed for magnesium dis-
solution from the silicate and magnesium precipitation as the carbonate, whereby rapid dissolution
requires a low pH of around 1 while rapid precipitation requires a considerably higher pH of around 8.
This paper investigates a novel concept which is to use a tertiary amine to bind with protons and raise
the pH to around 8 and to then regenerate the amine through the use of heat due to the strength of
the amine–proton bond decreasing with increasing temperature. This approach provides the low pH
and high temperature that is needed for Mg dissolution and the high pH need for carbonate precipitation.
The amine can be thought of as a regenerable buffer.
Dissolution of Mg from serpentine has been found to be favourable with a solids to solution volume of
more than 50 g/L to enable a low pH, and with temperatures close to the boiling point of the solution. The
pH needed for magnesium carbonate precipitation was found to be approximately 8.2. Both triethyl-
amine and tripropylamine were found to be capable of achieving this at 18 °C. Yields of around 20–
40 wt.% carbonate were achieved using residence times of approximately 1 h. The pH swing for the ter-
tiary amines was found to be approximately 2.5 pH units between 5 and 85 °C, suggesting that an amine
capable of achieving a pH of 8.2 at low temperature generates a pH of 5.7 in solution when heated to
85 °C. Further work will examine whether the lower pH values needed for serpentine dissolution can
be achieved by heating the protonated amine to higher temperatures.
Ó 2013 Elsevier Ltd. All rights reserved.
This paper investigates a novel concept to enable the conversion
of CO2 into mineral carbonates using an aqueous route. The barrier
that is currently stalling conversion is pH control. The ﬁrst part of
this introduction outlines why pH is critical, the second part out-
lines what work has been done in the ﬁeld of converting CO2 into
mineral carbonates using an aqueous route, and the third part
introduces the novel concept behind this study.
1.1. Thermodynamic modelling of the carbonate system
In order for a mineral carbonate, such as MgCO3, to form the
concentrations of CO2À
3 and Mg2+
in solution must be high enough
for Eq. (1) to be satisﬁed .
3 3:46 Â 10À8
The solubility of CO2 in aqueous solution and the dissociation of
carbonic acid to form bicarbonate and carbonate anions can be de-
scribed by Eqs. (2)–(5) [1–4].
ðatmÞ ¼ 29:36½H2CO3 ð2Þ
3 ¼ ½H2CO3=ð2:249 Â 106
0016-2361/$ - see front matter Ó 2013 Elsevier Ltd. All rights reserved.
⇑ Corresponding author. Tel.: +61 733653977; fax: +61 733654199.
E-mail address: firstname.lastname@example.org (K.M. Steel).
Fuel 111 (2013) 40–47
Contents lists available at SciVerse ScienceDirect
journal homepage: www.elsevier.com/locate/fuel
3 ¼ ½HCOÀ
3 =ð2:133 Â 1010
3 þ 2½CO2À
where 29.36 is the Henry’s law constant (dm3
) for CO2 in
water at 25 °C . Here, [H2CO3] includes both dissolved molecular
CO2 and molecular H2CO3. Eqs. (3) and (4) shows that CO2À
only start to become signiﬁcant in solution when the pH is above
about 8 (see Fig. 1). It could be thought that if the partial pressure
of CO2 was raised the concentration of all species in solution
3 and CO2À
3 ) would increase, and therefore a raised con-
centration of CO2À
3 would result. However, solving Eqs. (2)–(5) for
various partial pressures of CO2 (PPCO2
) shows that the concentra-
tion of CO2À
3 in solution stays very low at around 5 Â 10À11
spite the partial pressure of CO2 rising to 50 atm (see Fig. 2). This
is due to the H+
ions, forming from the dissociation of H2CO3, always
suppressing the formation of CO2À
3 (see Eqs. (3) and (4)). Only the
concentration of HCOÀ
3 becomes appreciable at high CO2 partial
Given the solubility limit of MgCl2 in solution (55 g/100 ml), the
maximum possible concentration of Mg2+
6 M Mg2+
and the corresponding carbonate concentration required
to precipitate MgCO3 is 5 Â 10À9
M. Given that 6 M Mg2+
to achieve as it is close to the solubility limit, a more realistic con-
centration of 0.6 M Mg2+
would require a CO2À
3 concentration of
5 Â 10À8
3 . Assuming a partial pressure of CO2 of 1 atm,
Eqs. (2)–(4) can be solved to establish a relationship between car-
bonate concentration and pH. Fig. 3 shows this relationship and
shows that when the carbonate concentration is 5 Â 10À8
equilibrium pH is approximately 5.5. This means that in order to
precipitate MgCO3 the pH needs to exceed 5.5.
1.2. Work to date on converting CO2 into mineral carbonates using the
Bond and co-workers [5–7] looked at the potential of using an
enzyme, called carbonic anhydrase, to convert power station CO2
into CaCO3. Coral reefs use carbonic anhydrase to assist in the pro-
duction of CO2À
3 . The enzyme was found to catalyse the formation
3 , however, although the catalyst enables the rapid forma-
tion of CO2À
3 , there is also a simultaneous rapid drop in solution
pH to approximately 4, which prevents the precipitation of carbon-
ates. In the work of Bond et al. [5–7] and in recent work by Mirja-
fari et al. , Ozdemir  and Rayalu and co-workers [10,11],
precipitation of carbonates was only possible if a buffer was added
to the solution in order to complex H+
ions and keep the pH high.
The buffer used was tris(hydroxymethyl)aminomethane or ‘Tris’,
which is commonly used to maintain pH at around 8. This buffer
could not be continuously used in a large scale process, as it cannot
be regenerated, which presents a major obstacle to the further
development of this approach to sequestering CO2. Coral reefs have
the ocean as a giant buffer for the H+
that they generate. It is
worthwhile to note that Mirjafari et al.  found that calcium car-
bonate did not precipitate when they used carbonic anhydrase
with no buffer, but found precipitation to take place when they
used the buffer with no carbonic anhydrase.
Instead of converting salt solutions into carbonates, researchers
have also looked at converting Mg–silicate minerals into carbon-
ates in order to use the neutralising capacity of the mineral .
O’Connor and co-workers [13,14] found they could convert 34%
of serpentine (Mg3Si2O5(OH)4) to MgCO3 using aqueous CO2 at
115 atm, 185 °C and a residence time of 24 h. The pH generated
in solution would be around 3. The high temperature would be
assisting the kinetics of Mg dissolution. Presumably the pH in-
creases as dissolution proceeds and this aids precipitation of car-
bonate. However, as the pH increases the dissolution rate of Mg
would drop to negligible levels. The silica left behind and carbon-
ate forming would also hinder further Mg dissolution/carbonation.
O’Connor et al. found that artiﬁcially adding NaHCO3 to shift the
equilibria in favour of a higher CO2À
3 to H+
ratio aided carbonation.
1 2 3 4 5 6 7 8 9 10 11 12 13 14
Fig. 1. Equilibrium distribution of H2CO3, HCOÀ
3 and CO2À
3 species in solution.
partial pressure (atm)
Fig. 2. Equilibrium concentration of H2CO3, HCOÀ
3 and CO2À
3 in water as a function
of CO2 partial pressure (derived from Eqs. (2)–(5)).
1 2 3 4 5 6 7 8 9 10 11
Fig. 3. Equilibrium concentration of CO2À
3 as a function of pH for CO2 partial
pressures of 0.1, 1, 10 and 100 atm.
K.M. Steel et al. / Fuel 111 (2013) 40–47 41
It follows that the best way to enable carbonation might be to
have a two stage process where the ﬁrst stage is optimised for
Mg dissolution (high temperature and low pH) and the second is
optimised for Mg carbonation (high pH). This need has been recog-
nised as pH swing . pH swing requires the removal of H+
from solution, which is usually achieved with soluble oxides/
hydroxides, however, these come from the calcination of carbon-
ates which is obviously not possible.
1.3. Novel concept for converting CO2 into mineral carbonates using
the aqueous route
This paper investigates the use of weakly basic tertiary amines
for complexing H+
ions generated when CO2 is added to solution
and therefore enabling carbonate precipitation from a variety of
salt solutions. The acid dissociation constant for protonated ter-
tiary amines varies with temperature and so the approach is to
use ‘‘temperature swing’’ to regenerate the amine, whereby the
loaded amine is heated to strip the acid off. This concept has re-
cently been patented .
Conventional CO2 capture technologies involve absorbing CO2
into a mixture of primary and tertiary amines, including monoeth-
anolamine (MEA) and methyldiethanolamine (MDEA) respectively.
The reason for this is that while the primary amine forms a strong
carbamate bond and therefore enables a high CO2 loading due to
the strength of the bond considerable energy is needed to break
it and regenerate the MEA (approximately 3–5 MJ/kg CO2 ).
In order to reduce the energy load tertiary amines are blended. Ter-
tiary amines do not bond with CO2 but rather strip the solution of
ions thus driving the formation of HCOÀ
3 and CO2À
3 ions in solu-
tion. The loading of CO2 in solution as these ions is much lower but
the energy needed to regenerate the MDEA is much less, such that
the total energy needed to regenerate the MEA/MDEA mixture is
around 1–3 MJ/kg CO2 . The introduction of MDEA means that
taller towers are needed for a given separation efﬁciency and so a
balance between capital and operating (energy) cost must be
The idea put forward here is to use tertiary amines alone to con-
vert CO2 into mineral carbonates. It is thought that the loading of
CO2 can be higher than that for conventional CO2 capture using
MDEA because the CO2À
3 ions that form leave the solution as solid
carbonate, thus providing a stronger driving force for the capture of
CO2 into solution.
The advantage of this concept for CO2 sequestration is twofold.
Firstly, the energy needed for the process might be considerably
lower than that needed for conventional CO2 capture because less
energy is needed to regenerate the amine and compression of the
CO2 (needed for storage) is not necessary. Secondly, the CO2 would
be locked up in a mineral form that is known to be stable for mil-
lions of years, which is an important consideration given the scale
of CO2 that needs to be stored.
Fig. 4 shows a simpliﬁed diagram of the novel concept that is
being explored. In stage 1, acid loaded amine is heated to
$100 °C and contacted with a Mg silicate such as serpentine
((Mg, Fe)3Si2O5(OH)4). At the high temperature the acid (HCl) dis-
sociates from the amine, thereby providing a low pH capable of
dissolving Mg out of the serpentine. The Mg depleted serpentine
is separated by density and/or ﬁltration and the solution, contain-
ing MgCl2 and regenerated amine, passes to stage 2. In stage 2, the
solution is cooled and the ﬂue gas containing CO2 is sparged
through. At the low temperature acid generated by the dissociation
of H2CO3 plus excess acid generated in stage 1 are complexed by
the tertiary amine which causes the pH to rise and consequently
3 concentration to rise to a level that is sufﬁcient to begin
interacting with Mg2+
and precipitating MgCO3. The MgCO3 is sep-
arated by density and/or ﬁltration and the solution, containing acid
loaded amine is recycled to stage 1.
The overall reaction taking place in stage 1 is as follows:
Mg3Si2O5ðOHÞ4 þ 6R3NHCl ! 3MgCl2 þ 6R3N þ 5H2O
þ 2SiO2 ðR1Þ
And the overall reaction taking place in stage 2 is as follows:
MgCl2 þ CO2 þ H2O þ 2R3N ¡ MgCO3 þ 2R3NHCl ðR2Þ
The concept shown in Fig. 4 can be used in a variety of different
ways by a variety of different industries, and is not locked into one
The concept could be used for the treatment of magnesium sil-
icate deposits as described above. It is worthwhile to note that
many magnesium silicate deposits contain signiﬁcant levels of
valuable heavy metals such as Ni, called Ni laterites, and the pro-
cess could therefore have the dual operation of Ni extraction com-
bined with CO2 sequestration. It is known that treating Ni laterites
with acid to dissolve the Mg enables a greater amount of Ni to be
extracted into solution.
Secondly, the concept could be used for salt mining operations.
Chloride and sulphate salt deposits are mined for KCl or K2SO4 to
be used as fertiliser. By using solution mining the salt comes up
hot and is cooled to separate the potassium salts. The refuse salt
represents a waste that is generally deposited back into the re-
serve. The salt solution could be processed using the sequestration
technology shown in Fig. 4 to form a mixture of carbonates and
bicarbonates that are deposited back into the reserve for longterm
CO2 storage. The heat that must be taken out of the solution as it
comes to the surface could be used to provide the heat needed
for stage 1. Using the concept for salt solutions means that a by-
product of the process is acid, either hydrochloric or sulphuric acid.
Therefore, the scale of the operation would need to be matched
with HCl or H2SO4 needs in the oil, chemical and mineral sectors.
This paper presents our work to date on this novel concept.
Fig. 4. Simpliﬁed block ﬂow diagram of the proposed CO2 sequestration technology.
42 K.M. Steel et al. / Fuel 111 (2013) 40–47
2.1. Serpentine sample
The serpentinite sample (Mg3Si2O5(OH)4) used for this study
was obtained from a naturally occurring deposit in northern
Queensland, Australia. It was initially ground by hand in a pestle
and mortar and then in a laboratory attrition mill. The ground ser-
pentinite was sieved with ASTM standard sieves to obtain particles
with a diameter of 57 lm. Australian Laboratory Services (ALSs)
performed the elemental analysis via alkali fusion, acid digestion
and inductively coupled plasma-atomic emission spectroscopy
(ICP-AES) of the resulting solution. The loss on ignition at
1000 °C (LOI1000) was also performed using a TGA furnace. The re-
sults of ALS’s analysis based on their method ME-ICP85 (Silicates
by Fusion, ICP-AES) and ME-GRA05 (H2O/LOI by TGA Furnace)
are summarised in Table 1.
Mineral composition was ﬁrst probed via X-ray diffraction anal-
ysis using a PANanalytical XPERT-PRO diffractometer with Cu Ka
target (k = 0.15406 nm) at room temperature. Measurements were
made in a step scan mode (0.1°/step) over the 2h range of 10–90°.
Phase matching of the X-ray powder diffraction (XRPD) pattern of
the serpentinite sample against the International Centre for Dif-
fraction Data (ICDD) database suggested antigorite to be the pri-
mary serpentine phase present. A calibrated TA Instrument
SDTQ600 Thermogravimetric analyser-differential scanning calo-
rimeter provided further mineral characterisation via thermogravi-
metry–derivative thermogravimetric analysis (TGA–DTG).
Replicate runs were obtained for 10 mg of the À53 lm samples
using alumina crucibles and heated from 30 °C to 1000 °C at a heat-
ing rate (b) of 10 °C minÀ1
. A 10-min isothermal stage was em-
ployed at 110 °C to determine the moisture content of the
material and was found to contain 1.0 wt.%. The total mass loss
of the dry sample (Dm105–850 °C) was determined as
11.5 ± 0.01 wt.%, which is in fair agreement with the analysis made
by ALS (11.7 wt.%).
Fig. 5 shows the TGA–DTG proﬁle of the sample, where the
characteristic serpentine doublet comprising the DTG temperature
shoulder, Tsh, showing at 597 °C and the peak temperature, Tp, at
718 °C. Thermal analysis suggests that this particular serpentinite
sample is fully serpentinized and contains antigorite as well as
lizardite (antigorite + lizardite). The antigorite component displays
its DTG peak temperature, Tp1ATG at 718 °C and its diagnostic peak,
Tp2ATG at 747 °C. The shoulder at 701 °C is thought to indicate the
lizardite component, Tp1LIZ. Based on XRPD and TGA-DTG analysis,
we then refer to this sample as ‘‘serpentinite’’, rather than antigor-
ite as this specimen also contains lizardite.
2.2. Magnesium dissolution
Magnesium dissolution experiments were carried out using AR
grade HCl (37 wt.%) and Millipore water. 0.5 g of sample and acid
solution was mixed in a 250 ml spherical ﬂat-bottom ﬂask
mounted on a magnetic stirrer/hotplate and equipped with a cold
water condenser. The effects of temperature, concentration of HCl,
residence time and acid solution volume were investigated. At the
duration of the experiment, the residue was vacuum ﬁltered, dried
overnight and weighed. The pH of the ﬁltrate solutions was mea-
sured using a pH electrode. The solid remaining was analysed by
ALS using the procedure described above to determine the extent
of Mg dissolution.
2.3. Carbonate precipitation
Carbonate precipitation experiments were carried out using a
250 ml Erlenmeyer ﬂask open to the atmosphere. Food grade CO2
was injected into the solution through a sparger containing ﬁve
holes of 2 mm diameter each. Pressure was regulated at 1.4 bar
and ﬂow was set at approximately 1 L/min using a rotameter.
While sparging the solution with CO2, tertiary amine was added
dropwise via a burette while simultaneously measuring pH via a
pH electrode. The amines investigated were simple straight chain
trialkylamines with increasing chain length, i.e. triethylamine, tri-
propylamine, tributylamine, tripentylamine, trihexylamine. The
solution was observed for the onset of precipitation. If a precipitate
form, it was ﬁltered through Whatman No. 1 ﬁlter paper, dried and
weighed. Precipitated solids were analysed by XRD and ICP-AES for
compound and elemental determinations, as described above.
2.4. Amine regeneration
The ability of the amines to be regenerated and liberate bound
acid was investigated via a series of titrations at various tempera-
tures whereby the amines were added to a standardised solution of
0.1 M HCl (10 ml). The HCl was placed in a 3-neck round bottom
ﬂask which was immersed in a water bath to control temperature.
A condenser was ﬁtted vertically to the middle neck and a ther-
mometer and pH electrode were inserted and sealed through each
of the side necks. Amine was added through the opening at the top
of the condenser using an automatic pipette. For experiments per-
formed at 5 °C ice was added to the water bath. The tertiary amines
studied were triethylamine, tripropylamine, tributylamine and
3. Results and discussion
3.1. Magnesium dissolution
The effect of HCl concentration on the dissolution of Mg from
serpentine is shown in Fig. 6. All percentages are weight percent-
ages. The residence time used was 3 h and the temperature was
the boiling temperature of the solution ($100 °C). The stoichiome-
tric amount of acid needed to dissolve all of the Mg according to R1
is approximately 0.12 M which gives approximately 40% extrac-
tion, while a plateau of approximately 65% extraction is reached
at around 0.5 M HCl. Fe and Al were also found to dissolve with
similar extraction levels to those of Mg, showing that the elements
do not appear to dissolve selectively with respect to HCl
Because it is desirable to not have excess acid in solution to
minimise the energy needed for amine regeneration a compromise
between Mg extraction and solution pH must be struck. Fig. 6
shows the pH change as a function of HCl concentration. With
twice the stoichiometric amount needed, the pH of the spent solu-
tion is still $1.
The effect of residence time on dissolution of Mg in 0.25 M HCl
is shown in Fig. 7. Equilibrium is reached after approximately 3 h.
The effect of temperature on the dissolution of Mg in 0.25 M HCl is
shown in Fig. 8. For temperatures less than 50 °C only a small
amount of Mg dissolves. At temperatures higher than 50 °C a linear
Chemical composition of the serpentinite sample.
39.3 44.1 7.38 1.03 0.35 0.24 0.10 0.04 11.7
Values obtained via inductively coupled plasma-atomic emission spectroscopy on
fused samples after acid digestion. Based on ALS’s ME-ICP85.
Value obtained by heating the moisture free sample to 1000 °C using TGA fur-
nace. Based on ALS’s ME-GRA05.
K.M. Steel et al. / Fuel 111 (2013) 40–47 43
increase in Mg dissolution with respect to temperature is obtained,
reaching approximately 65% at $100 °C and suggesting that at
temperatures above 100 °C higher extraction efﬁciencies might
be achieved. Extrapolation suggests that close to 100% extraction
might be achieved at 140 °C. Experiments using a pressurised ves-
sel to enable higher temperatures above 100 °C are recommended
to conﬁrm the extrapolated trend shown.
As presented in the introduction, it is desirable to have a high
concentration in solution as this reduces the concentration
3 needed for MgCO3 precipitation. Experiments were per-
Fig. 5. The TGA–DTG proﬁle of the serpentinite sample used in this study with the characteristic peaks indicated.
Fig. 6. Effect of HCl concentration on the dissolution of Mg and ﬁnal pH (residence
time 3 h, $100 °C, 0.5 g, 100 ml).
Fig. 7. Effect of residence time on the dissolution of Mg (0.25 M HCl, solution and
$100 °C, 0.5 g, 100 ml).
Fig. 8. Effect of temperature on the dissolution of Mg (0.25 M HCl, residence time
3 h, 0.5 g, 100 ml).
Fig. 9. Effect of solution volume on the dissolution of Mg (0.025 mols HCl, residence
time 3 h, 0.5 g, $100 °C).
44 K.M. Steel et al. / Fuel 111 (2013) 40–47
formed keeping the amount of HCl the same and decreasing the
solution volume from 100 ml. Twice the stoichiometric amount
needed (0.024 mols) was chosen for the amount of HCl. Fig. 9
shows the effect of reducing the solution volume down to 10 ml.
The extraction increases as the solution volume decreases reaching
approximately 85% with only 10 ml of solution. The concentration
of Mg in solution is approximately 0.35 M.
This work has shown the importance of both acid concentration
and temperature on the dissolution of Mg. It is recommended to
operate with a solids to solution volume of more than 50 g/L for
the extraction stage, a temperature close to the boiling tempera-
ture of the solution or higher if using a pressurised vessel, a resi-
dence time of 3 h and concentration no more than twice the
stoichiometric amount needed for reaction. These conditions en-
able high extractions of Mg approaching 100%.
3.2. Carbonate precipitation
The extract solution from the experiment with 10 ml solution
volume shown in Fig. 9 was used for carbonation. Tripropylamine
was added dropwise. As the pH increased to approximately 5 a
light brown precipitate formed which was found from elemental
analysis to contain around 20.1 wt.% Fe, 10.8 wt.% Si and 5.5 wt.%
Al and only 0.2 wt.% Mg. This product comes from the hydrolysis
of Fe, Si and Al which dissolved during serpentine dissolution.
The leaching studies had shown that the elements dissolved simul-
taneously with Mg. After removing this precipitate further addi-
tions of TPA raised the pH to approximately 8 at which point a
white precipitate formed. This precipitate began forming within
a few minutes. After bubbling CO2 for 45 min, the precipitate was
recovered by ﬁltration, dried and analysed. The weight was
0.20 g. XRD analysis indicated the formation of nesquehonite
(MgCO3Á3H2O). Elemental analysis showed that the purity was
high with a composition of 18.8 wt.% Mg, 0.2 wt.% Ca, 0.06 wt.%
Fe and 0.01 wt.% each of Al and Si. The yield was 29.1 wt.% (i.e.
29.1 wt.% of the Mg extracted into solution was converted to the
MgCO3 precipitate). The remaining Mg would be recycled to the
ﬁrst stage of the process.
In order to study MgCO3 more precisely and the pH level needed
for carbonate precipitation without the hindrance of other dis-
solved elements model compound work was performed using
Mg(OH)2. 0.126 g of Mg(OH)2 was dissolved in 50 ml of 0.0965 M
HCl such that the acid was slightly in excess and gave a ﬁnal pH
of 2.08 after the Mg(OH)2 had completely dissolved. To this solu-
tion 5 Â 10À4
mols of TPA was added, which is the amount needed
to neutralise the excess acid. The pH rose to 9.59. CO2 was then
bubbled through the solution and after 5 min the pH had decreased
and stabilised at 4.62 and no precipitate had formed. A further
addition of TPA was made (0.006 mols) and the pH stabilised at
approximately 8.27. This addition of TPA is 1.5 times that needed
for reaction 2. Over the next half an hour, CO2 was continually bub-
bled through the solution. It was found that TPA need to be contin-
ually added dropwise in order to maintain the pH at a level above
8. The total amount of TPA added (neglecting the initial 5 Â 10À4
mols) was 0.021 mols and the ﬁnal pH was 8.43. The solution was
ﬁltered and the solid recovered and air dried. The weight of the so-
lid was 0.051 g and as found above, analysis indicated nesqueho-
nite (MgCO3Á3H2O) to be the primary phase present. The yield
was approximately 17% and the amine used was 5 times in excess.
The above experiment was repeated with slower additions of
TPA. A total amount of 0.0085 mols (twice excess) was added over
a period of 1 h. The mass of solid recovered was 0.073 g (dried)
which is a yield of 24 wt.%.
The above experiment was repeated with triethylamine (TEA).
0.164 g of Mg(OH)2 was dissolved in 50 ml of 0.112 M HCl, which
is the stoichiometric amount needed for complete dissolution.
CO2 was bubbled which decreased the pH to 5.38. 0.0072 mols of
TEA was added and the pH increased to 10.16 and simultaneously
the solution became milky with precipitation. As CO2 addition con-
tinued the pH decreased to 7.27 even though the amount of TEA
added was 30% above that needed for reaction 2. A further addition
of 0.0036 mols of TEA increased the pH to 9.32 initially but then it
stabilised at 8.23. The solution was ﬁltered to recover the solid,
which had a mass of 0.133 g and therefore a yield of 34 wt.%.
Tests with both tributylamine and tripentylamine did not yield
precipitates, which is thought to be due to the pH generated by the
amines not being high enough. It is possible that decreasing the
temperature would enable precipitation to take place with these
These experiments have shown that the pH needed for magne-
sium carbonate precipitation is approximately 8.2 and that trieth-
ylamine and tripropylamine are capable of achieving this. An
excess of amine has been found to be necessary to maintain the
pH of 8 while CO2 is bubbled through the solution. So far yields
of around 20–40 wt.% have been achieved. The reason why precip-
itation did not occur at lower pH levels, such as the pH level of 5.5
predicted from the theoretical modelling work reported in the
introduction, is thought to be due to the kinetics being too slow be-
low 8.2. It was found with TEA that precipitation was within sec-
onds when the pH was 10.
Further experiments will investigate the kinetics of carbonate
precipitation by sampling periodically, particularly during the
3.3. Amine regeneration
Based on the serpentine dissolution work, to achieve dissolu-
tion of Mg under reasonable conditions of a residence time of less
than an hour and reaction time of 100–150 °C, the pH of the solu-
tion needs to be approximately less than 1. Based on the carbon-
ation work, the pH of the Mg rich solution needs to be raised to
approximately 8.2. To examine the ability of tertiary amines to
reversibly enable this change a series of titrations were performed.
The dissociation constant for triethylamine at various tempera-
tures has been published by Hamborg and Versteeg , whereby
the pKa is 10.89 at 18 °C and decreases to about 9.17 at 90 °C.
Fig. 10 shows these constants converted into a titration curve
whereby the amine is being added to 0.1 M HCl. Our own titration
points are also shown for comparison. The titration curve is ex-
pressed this way as it mimics the real system. Total amine essen-
Fig. 10. Titration curve for triethylamine against 0.1 M HCl at 18 °C and 90 °C
generated from pKa data obtained from literature and at 18 °C generated from
K.M. Steel et al. / Fuel 111 (2013) 40–47 45
tially means the amount of amine added expressed as a concentra-
tion. At 20 °C the pH rises to approximately 11 while at 90 °C it
rises to approximately 9. While a pH of 11 would assist with the
precipitation of carbonates the pH of 9 at the higher temperature
will not assist with the dissolution of Mg from serpentine.
Figs. 11–13 shows our own titration curves for tripropylamine,
tributylamine and tripentylamine at various temperatures. For tri-
propylamine, the ﬁnal pH is approximately 9.5 at 18 °C and 7.1 at
85 °C. For tributylamine, the ﬁnal pH is approximately 8.6 at 5 °C
and 6.0 at 85 °C. For tripentylamine, the ﬁnal pH is approximately
6.5 at 5 °C and 4.0 at 85 °C. It follows that with a temperature rise
from 5 to 85 °C, the change in ﬁnal pH is approximately 2.5 pH
units. This work shows that at elevated temperatures an acid
loaded alkylamine with a long chain length behaves similarly to
a weak acid and therefore might have the ability to dissolve Mg
The equivalence point (point of highest gradient) was found for
each titration curve from which pKa values were estimated. Table 2
shows results from this analysis and Fig. 14 shows the constants
for tributylamine and tripentylamine as a function of temperature.
Lines of best ﬁt have been drawn through the data and extended to
135 °C. Fig. 14 also shows a line corresponding to the pKa values
reported by Hamborg and Versteeg  for triethylamine. The
lines obtained from this work appear to decrease more steeply
than those for triethylamine, which may be due to experimental
error associated with vapour losses from the system which could
concentrate the protons and give lower pH values.
Further experiments to obtain more accurate titration data
including data at higher temperatures and pressures (100–150 °C
and 1–5 bar) is planned for the future. The experiments at higher
temperature will also involve treating serpentine with the regener-
ated amine and acid mix to study the dissolution behaviour. These
experiments are akin to treating serpentine with weak organic
acids. There is currently a lack of studies on the behaviour of ser-
pentine with weak organic acids particularly at high temperatures
where the kinetics of dissolution is favourable. Teir et al.  have
reported the behaviour of carboxylic acids alongside stronger acids
however the studies were conﬁned to 20 °C. Unlike strong acids
which provide a high initial concentration of protons which de-
creases as the mineral dissolves, weak acids provide a low concen-
tration that is maintained as the mineral dissolves, and protons are
consumed in the acid-base reaction thereby driving the dissocia-
tion of more protons from the weak acid.
Fig. 11. Titration curve for tripropylamine against 0.1 M HCl at 18 and 85 °C.
Fig. 12. Titration curve for tributylamine against 0.1 M HCl at 5, 18 and 85 °C.
Fig. 13. Titration curve for tripentylamine against 0.1 M HCl at 5, 18 and 85 °C.
Calculated pKa values for tertiary amines at 5, 18 and 85 °C.
5 °C 18 °C 85 °C
Tripropylamine nd 9.47 7.32
Tributylamine 9.84 8.44 6.32
Tripentylamine 6.81 5.80 3.98
nd: Not determined.
Fig. 14. Expected trend for pKa as a function of temperature for tributylamine and
tripentylamine with lines of best ﬁt, and pKa line for triethylamine as derived from
Hamborg and Versteeg .
46 K.M. Steel et al. / Fuel 111 (2013) 40–47
If the regenerated amines enable high degrees of serpentine dis-
solution, experiments will move to examining the extent that the
amines can be continually recycled.
The best conditions for the dissolution of Mg from serpentine
have been found to be a solids to solution volume of more than
50 g/L to enable a high proton concentration. The amount of acid
should be no more than twice the stoichiometric amount needed
for reaction. Reaction temperature should be as high as possible,
close to the boiling temperature of the solution or higher (100–
150 °C) if using a pressurised vessel. These conditions combined
with a residence time of 3 h are able to dissolve approximately
85% of the Mg in serpentine.
These experiments have shown that the pH needed for magne-
sium carbonate precipitation is approximately 8.2 and that trieth-
ylamine and tripropylamine are capable of enabling this at 18 °C. It
appears that an excess of amine is needed to maintain the pH of 8
while CO2 is bubbled through the solution. So far yields of around
20–40 wt.% have been achieved for tripropylamine using residence
times of approximately 1 h. Precipitation occurred more rapidly for
triethylamine owing to the higher pH generated.
The association of tertiary amines with HCl has been found to
decrease with increasing temperature such that there is a differ-
ence of approximately 2.5 pH units between 5 and 85 °C. This
means that an amine capable of achieving a pH of 8.2 at low tem-
perature generates a pH of 5.7 in solution when heated to 85 °C.
While this is not low enough to provide a high rate of serpentine
dissolution it is thought that increasing the temperature beyond
85 °C may yield pH levels capable of dissolving high levels of Mg,
particularly given that high temperatures aid the kinetics of disso-
lution. Further work is required to study the dissolution behaviour
of serpentine with regenerated amine solutions at elevated tem-
peratures and pressures.
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