International Flame Research Foundation
Finnish-Swedish Flame Days January 28 -29, 2009
Naantali, Finland
CO2 EMISSIONS MITIGATION IN FINLAND BY CARBONATION
OF SILICATE MINERALS AND INDUSTRIAL WASTES
R. ZEVENHOVEN*, J. FAGERLUND, A. WIKLUND, E. NDUAGU,
A.-M. FORSMAN, M. FÄLT, H.-P. MATTILA
Åbo Akademi University (ÅA), Heat Engineering Laboratory
Biskopsgatan 8, FI-20500 Åbo/Turku Finland
*Corresponding author: ron.zevenhoven@abo.fi, +358 2 2153223
S. ELONEVA, A. SAID
Helsinki University of Technology (TKK), Department of Energy Technology
PO Box 4400, FI-02015 Espoo Finland
ABSTRACT: Carbon dioxide capture and storage (CCS) technology has to be
implemented also in Finland if fossil, carbon-containing fuels are to be used in the
future. As underground cavities for geological storage of CO2 appear to be absent in
Finland the CO2-producing energy sector seems to aim at carbon dioxide capture and
export (CCE) instead. However, as recognised primarily by Finnish pulp and paper
industry and metal industries, Finland can make use of its significant resources of
magnesium silicates and industrial wastes. Finland’s CO2 emissions in excess of the
1997 Kyoto Protocol commitments are currently 10 – 15 Mt/a (megatons).
The carbonation of Finnish magnesium (Mg) silicate resources appears to offer enough
capacity for storing at least several Gt (gigatons) CO2, forming a benign carbonate
product that doesn’t require post-storage monitoring of the CO2. Besides this, the
carbonation of slag by-products from iron- and steelmaking allows for producing many
kt (kilotons) of valuable precipitated calcium (Ca) carbonate (PCC) products with
applications in paper-making, or for flue gas desulphurisation.
Work on CO2 mineralisation has been ongoing at TKK and ÅA for many years, with
increasing support from industrial partners from Finland and abroad. In this paper,
results and achievements so far are summarised.
For long-term storage of large amounts of CO2, a route is being optimised that involves
the production of Mg-hydroxide from serpentinite rock material, followed by
pressurised fluidised bed (PFB) carbonation to Mg-carbonate. The last (i.e. carbonation)
step should give the energy release that must cover for the process’ energy input needs.
For the iron-and steel sector, slag by-products are processed in a wet, aqueous process
route that involves extraction of calcium followed by carbonation with CO2 and
recovery of the solvent chemicals. The gas used may be a flue gas after removal of fly
ash and possibly also nitrogen- and sulphur oxides and some trace elements.
Process routes and the quality and amounts of carbonate products shall be presented.
Keywords: CO2 mitigation, mineralisation, magnesium carbonate, calcium carbonate
1.
INTRODUCTION
Carbon dioxide capture and storage (CCS) is an important option for mitigating CO2
emissions from human activities (IPCC, 2005). It comprises separation and compression
of CO2 from industry and power plants and transportation of CO2 to a storage site
followed by long-term storage. The carbonation of natural silicate minerals that contain
alkaline-earth oxides like magnesium oxide (MgO) and calcium oxide (CaO) offers a
leakage-free alternative to using underground geological formations for the long-term
storage of CO2. One of the countries that can make use of this is Finland, where it was
assessed that underground storage sites are not available while vast resources of suitable
magnesium silicates exist (Koljonen et al., 2004). Finland has emission reduction
commitments under the 1997 Kyoto protocol to reduce its greenhouse gas emissions to
the level of year 1990 during 2008-2012, while current emissions of CO2 are 10-15 Mt/a
above that level. For example, the serpentinite rock formations in Eastern-Central
Finland alone could bind several Gt of CO2 (Teir, 2008, Rinne, 2008).
Naturally occurring silicate minerals such as serpentine and olivine contain high
concentrations of MgO, while pyroxenes and amphiboles are a potential source for both
CaO and MgO. Carbonation of these traps CO2 as environmentally stable solid
carbonates, which for serpentine (R1) and wollastonite (R2) minerals can be described
by the following overall chemistry:
Mg3Si2O5(OH)4 (s) + 3CO2 (g) → 3MgCO3 (s) + 2SiO2 (s) + 2H2O (l)
(R1)
CaSiO3 (s) + CO2 (g) → CaCO3 (s) + SiO2 (s)
(R2)
Also Mg- or Ca- containing mine tailings and by-products or wastes from industry may
be used – see Zevenhoven et al. 2006b, Eloneva 2008, Teir 2008. Only the carbonation
of Mg-based material offers the capacity needed for significantly reducing
anthropogenic CO2 emissions and this is covered in the first part of this paper. In the
second part the production of valuable precipitated calcium carbonate (PCC) from slag
by-products from iron- and steelmaking is addressed; this industry sector, the mineral
processing sector and the pulp and paper sector have started considering CO2
mineralisation as a route to CO2 emissions mitigation, a route to increases profits from
by-products and/or a cheap route to commercial carbonate materials. Unfortunately,
Finland’s energy sector hasn’t yet embraced CO2 mineralisation as the CCS-route to
follow and appears to prefer carbon dioxide capture and export, CCE (T&T, 2008).
Moreover, a significant renewed effort was started to locate geological sites inside
Finland’s borders for carbon dioxide capture and geological storage (further referred to
as CCGS1) and according to recent reporting these have not yet been found (Nieminen,
2008). Mapping studies for CCGS “(GeoCapacity”) at the European level are ongoing
(Vangkilde-Pedersen et al., 2008), unfortunately Finland’s data is not included here.
1
The recent proposal for an EU directive on geological storage of carbon dioxide (CEC, 2008) confuses
this CCS method with CCS in general: “...capture CO2 from industrial installations and store it in
geological formations (carbon dioxide capture and storage, or CCS)”.
2/15
This paper is an updated compilation of two papers presented by Zevenhoven and
Eloneva at ACEME-08, Rome, October 1-3, 2008 (Zevenhoven et al., 2008b; Eloneva
et al., 2008c).
Ocean carbon
Biomass
Atmosphere
Annual
emissions
CARBONATION OF MAGNESIUM SILICATE MINERAL FOR LONGTERM STORAGE OF CARBON DIOXIDE
Mineral
Carbonates
100 000
Underground
injection
10 000
EOR
1 000
Fossil carbon
Characteristic storage time (yr)
2.
100
10
1
1
10
100
1 000
10 000 100 000 1 000 000
Carbon storage capacity (Gt)
Figure 1. Estimated storage capacities and storage times for various CCS methods.
(from Zevenhoven et al., 2006b, after Lackner, 2003)
As Figure 1 illustrates, (Mg-based) mineral carbonation gives the highest capacity and
longest storage time of all currently known CCS options (Lackner, 2003), and in
contrast to CCGS, post-storage monitoring of CO2 will not be needed. Mineralisation of
CO2 is attractive especially at locations where absence of underground formations
excludes CCGS, the risk of leakage of underground stored CO2 is considered
unacceptable, or large resources of material suitable for carbonation are present. Besides
this, the deployment of CCGS is progressing rather slowly (current total capacity ~7 Mt
CO2/a) and might not offer the necessary capacity soon enough, putting pressure on the
development of alternatives.
However, no commercial silicate mineral carbonation technology yet exists. Although
the overall reactions (R1) and (R2) are exothermic (which implies that proper
optimisation and process integration should allow for operation at zero or negative net
energy input) the natural carbonation of silicate minerals is very slow. The reaction
must be accelerated considerably to be an economically viable large-scale CO2 storage
technology. Technology overviews were given by Newall et al. (2000), Huijgen and
Comans (2003, 2005), and recently by Sipilä et al. (2008), dealing with Mg-based as
well as Ca-based materials. Two problems must be solved to make large-scale
mineralisation of CO2 more attractive (Herzog, 2002):
•
•
extracting or activating the reactive component MgO from silicate mineral, and
speeding-up the carbonation chemistry kinetics
3/15
A small yet steadily increasing number of countries are looking into technical solutions
to perform the carbonation of magnesium with CO2 on a large scale; the most important
results were reported from the U.S. since the mid 1990’s. While most research
concentrates increasingly on methods using aqueous solutions, research in Finland still
addresses gas/solid methods. The reason for preferring a gas/solid route is that aqueous
processes have temperatures much lower than, say, 500°C where the rate of the
gas/solid carbonation with MgO or Mg(OH)2 appears to become significant. The energy
consumption of the “wet” processes is high; it should be much lower (or negative) for a
gas/solid process with heat recovery. An important feature of CO2 mineralisation will be
solids mining and handling and the question of what to do with the carbonate product
material. A metric ton of CO2 will require 2.5-3 tons of magnesium silicate mineral (for
coal combustion-derived CO2 this implies ~8 tons of mineral per ton coal). Altogether
this will result in a mining activity similar to typical commercial mining of coal or
metal-containing ore.
In the open literature, hardly any data can be found on the kinetics of chemical reactions
between magnesium silicates or oxides with CO2. Since the mid-1990’s many results
(mainly qualitative work-in-progress reporting) from mineral carbonation work by U.S.
consortia were published. Roughly after year 2000 countries like Finland joined in.
2.1
Gas-phase conversion routes
The first experiments reported from the U.S. during the 1990’s showed low levels of
carbonation of magnesium silicates even at elevated (supercritical) CO2 pressures and
temperatures like 340 bar, 500°C for sub-mm size particles after several hours. This and
the results from later work in the U.S. and Finland are summarised in Table 1. Recent
tests made in Finland with a serpentinite mine tailing sample (~83 %-wt serpentine +
~14 %-wt magnetite Fe3O4 + 3 %-wt others, size < 75 µm) from the Hitura nickel mine
in central Finland with (sub- and) supercritical CO2 involved direct carbonation during
3½h after heat-up and pressurisation to 57-300 bar, 150-200°C gave < 2% carbonation.
After efforts to directly carbonate Mg-silicates the U.S. work shifted to MgO extraction
from magnesium silicates and the effect of water (released from Mg(OH)2, for example)
on carbonation. The carbonation of Mg(OH)2 was soon found to be significantly faster
than carbonation of MgO, and the effect of pressure (of H2O and CO2) on carbonation
was studied in detail. Increasing pressures above the minimum CO2 pressure for stable
MgCO3, slows down the dehydroxylation of Mg(OH)2, generating fewer reactive MgO
sites for carbonation (Butt et al., 1996, Béarat et al., 2002, Chizmeshya et al., 2002).
Attempts to produce Mg(OH)2 from Mg-silicates resulted in the evaluation of a process
using hydrochloric acid (HCl in water) with intermediates MgCl2·nH2O and Mg(OH)Cl,
followed by gas/solid carbonation of Mg(OH)2 (see Newall et al., 2000). Due to the
complexity and energy consumption of this process route other U.S. researchers
proceeded with working on direct routes based on aqueous solutions (see below).
4/15
Table 1. Magnesium-silicate mineral carbonation since the 1990s: gas/solid
Mineral
Temp.
ºC
Press.
bar
Part. size
µm
Time
h
Conversion %
Where,
when
Serpentinite
500
340
50-100
2
25-30
US 1997
1, 20
~ 50
3
negligible
Fin 2000
Serpentinite
200,
1000→200
Serpentinite
150-200
57-300
75-125
3.5
<2
Fin 2006
Mg(OH)2
565
52
20
0.5
90
US 1997
Mg(OH)2
375
52
50-100
12
16.5
US 1996
MgO from
Mg(OH)2
1000→200
1, 20
~ 20
3
<5
Fin 2001 2003
Mg(OH)2
370 - 540
1-45
75-125
6
< 60
Fin 2004 2006
200,
In Finland, stepwise carbonation initially involved testing with Finnish serpentine and
(calcined) magnesium hydroxide Mg(OH)2 powder in a pressurised thermogravimetric
analyser (PTGA). Since elevated pressures didn’t give an increase in MgO carbonation
rate it was decided to proceed with Mg(OH)2 under test conditions where MgO
formation is thermodynamically unfavourable. (Zevenhoven et al., 2002, Zevenhoven
and Teir, 2004). Besides the PTGA tests, tests with Mg(OH)2 and MgO in an
atmospheric bubbling fluidised bed (FB) reactor were made, with CO2 as fluidising gas.
The product of carbonate material that builds up on the reacting particles (eventually
slowing down conversion) was noticeably removed from the particles as fines by
attrition and abrasion, which were entrained from the reactor with the exit gas flow.
These had a considerably higher MgCO3 content (8.1 %-wt) than the material in the bed
(4.4 %-wt) after ~11 h. (Zevenhoven and Teir, 2004). Clearly, FB reactors are very
suitable here.
PTGA tests on carbonation of Mg(OH)2 (97%, 75-125 µm) proceeded with tests at
pressures up to 45 bar. These confirmed that Mg(OH)2 can be carbonated faster than
MgO. Pressurisation allows for using higher temperatures, but for a given pressure the
carbonation rate decreases with temperature as a result of thermodynamic limitations –
see Figure 2a. XRD analysis confirmed a significant MgCO3 content in the end product.
Conversion at 40 and 45 bar was much faster than at 35 bar at similar or slightly higher
temperatures but was slow compared to the aqueous processes used in the U.S. (see
below) despite the progress made – see Figure 2b. Unreacted shrinking core (USC)
modelling was used to determine rate-limiting steps and rate parameters, showing that
after ~5% carbonation, so-called product layer diffusion becomes rate-determining.
(Zevenhoven et al., 2006a). This again illustrates that it would be beneficial to use a
fluidised bed reactor.
5/15
1000
100
50
10
40
1
Years
Carbonation efficiency (%)
60
63
s
ar
ye
3
ys
da
0
20
0.1
30
0.01
20
0.001
s
ur
ho
s
s
ur
ur
ho
ho
5
6
4<
0.0001
10
2002 2003
0
370
460
495
510
Tempera
ture (°C
)
525
2006
h
?
45
ar)
b
12
re (
u
1
s
s
Pre
35
540
2004 2005
Year
1
40
Figure 2a (left) Results of conversion of Mg(OH)2 (75-125 µm) to MgCO3 after 6 h for
various temperature/pressure combinations in 99%/1% CO2/H2O (Zevenhoven et al.,
2006a)
Figure 2b (right) Progress made in Finland as calculated reaction times for full
carbonation of Mg(OH)2 under chemical kinetics control (Sipilä et al., 2008)
The last PTGA test series made aimed at carbonation of serpentinite rock from the
Hitura nickel mine mentioned above (75-125µm) in three steps, producing MgO,
hydroxylation of MgO to Mg(OH)2 and carbonation of Mg(OH)2, respectively. It
showed that heat-up of serpentinite doesn’t produce MgO but converts serpentine into
olivine, water and quartz. Besides this, the dehydroxylation/ hydroxylation of Mg(OH)2
was studied in separate tests, confirming that conversion of MgO to Mg(OH)2 in
pressurised steam is very slow. The production of Mg(OH)2 presents yet another
challenge.
2.2
Aqueous phase conversion routes
The problems with the U.S. route using MgCl2 as intermediate resulted in the
development of a simpler, direct carbonation process using aqueous solutions. Soon
after the end of the 1990’s U.S. researchers reported conversion rates such as 65%
conversion after 1 h and 80% conversion within ½h after optimisation of solution
chemistry, heat treatment and grinding. Currently this method is considered to be the
most successful route for mineral carbonation using an aqueous solution of 0.64 M
NaHCO3 and 1 M NaCl at 185°C/150 bar (olivine), 155°C/115 bar (heat treated
serpentine) or 40 bar/100°C (wollastonite), respectively. The minimum costs for CO2
sequestration would be 54, 78 and 64 US$/ton CO2, respectively. For serpentine,
thermal pre-treatment (at 615-630°C) was found more efficient than mechanical
treatment (O’Connor et al., 2005, Gerdemann, et al. 2007). Herzog (2002) concludes in
his assessment that the reported carbonation levels and rates are obtained only at high
energy cost: “a 20% energy penalty for a coal-fired power plant”. Cost levels mentioned
6/15
make this method too expensive at this point – see also (IPCC, 2005). Nonetheless this
route still receives attention and improvements are being reported. It has been noted
recently, however, that the process costs are overestimated partly due to unrealistic
calculation of energy efficiency and energy costs: the costs of process heat input are
significantly over-estimated when charged the same way as power input, giving a false
impression of overall process economics. (Sipilä et al., 2008, Zevenhoven et al., 2008b).
The extensive pre-treatment of feedstock mineral makes the aqueous processes using
direct carbonation in fact indirect. As an alternative to energy-intensive pre-treatment
research teams worldwide embarked on methods to dissolve or leach Mg or Ca from
minerals or industial by-products and wastes using strong or weak acids, alkali solutions
or ligands. (See Huijgen and Comans 2003, 2005, Sipilä et al. 2008).
For example, Teir and co-workers (2007a,b) dissolved serpentinite from the Hitura
nickel mine mentioned above in both weak and strong aqueous solutions of common
acids or bases at room temperature during 1 h. It was found that H2SO4 extracted most
Mg, followed by HCl, HNO3, HCOOH and CH3COOH. None of the acids extracted Mg
selectively, extracting also some Fe and Si. Carbonate precipitation experiments were
made with two solutions prepared from serpentinite in strong solutions of HNO3 and
HCl, respectively; after filtration and drying, the remaining salts were dissolved in water
and iron oxides were precipitated. The magnesium-rich solutions (mainly dissolved
magnesium nitrates or chlorides) were used to precipitate magnesium carbonates by
exposing each solution to a CO2 flow while regulating pH using aqueous NaOH which
was necessary for producing precipitates. Quite pure hydro-magnesite,
Mg5(CO3)4(OH)2·4H2O was produced. However, the costs for pH regulating agent
NaOH and for make-up acid make this route economically unviable.
2.3
Magnesium carbonate product stability and use
Although large-scale mineralisation of CO2 allows for compact storage compared to
storage as (super-critical) fluid, the surface disposal of the carbonate material may be
problematic and unwanted at certain locations. Using MgCO3-containing material in, for
example, building materials or flame retardants may be considered; in paper products
however it can’t be used (Zevenhoven et al., 2006b). On the other hand, being also a
product of natural weathering of silicate mineral, the carbonate is benign and stable
when disposed of in nature. Tests showed that MgCO3 and CaCO3 are stable in rain
water and acid solutions down to pH 1 (Teir et al., 2006), but a very slow dissolving in
water bodies and streams could in fact be beneficial2.
Increasing the concentrations of Mg2+, Ca2+ and HCO3- ions can provide a buffer to the decomposition
of solid carbonate material under water, as the equilibrium reactions CaCO3(s) + H2O(l) + CO2(aq) =
Ca2+(aq) + CO32-(aq) + H2O(l) + CO2(aq) = Ca2+(aq) + 2HCO3-(aq), with HCO3- being much better
soluble than CO32-, will then be pushed to the left side. Thus preferable disposal sites for large amounts of
MgCO3 product could be the open sea or ocean.
2
7/15
2.4
Current research in Finland
Current work on carbonation of magnesium silicates in Finland focuses primarily on the
use of gas/solid fluidised bed (FB) reactors operating at carbonation temperatures and
pressures up to 600°C, 100 bar (super-critical conditions for CO2), supported by earlier
experiments using (PTGA) and extensive thermodynamic analysis (e.g. Zevenhoven et
al., 2006a, 2008). A test facility for this – see Figure 3 – was built and the
experimenting has started (Fagerlund et al., 2008). One part of the work involves
carbonation of Mg(OH)2 and MgO while at the same time developing methods (aqueous
or gas/solid) for producing this from serpentinite rock (Nduagu, 2008). The process as it
is currently being optimised at Åbo Akademi University and is shown schematically in
Figure 4.
Figure 3. Schematic lay-out of the fluidised bed set-up at Åbo Akademi University,
showing the (vertical) fluidised bed reactor (i.d. ~1.4 cm, height ~40 cm) and the
(horizontal) coil-shape gas pre-heating. In the far right upper corner the cyclone for
separating particulates from the gas stream.
CO2
Water
Ammonia
Ammonia
Other gas
Serpentinite
Mg(OH)2
Ammonium sulphate
MgCO3
Water
Iron
Ammonium sulphate
Sillica etc.
Figure 4. Schematic of the staged route for magnesium silicate carbonation
8/15
3.
CARBONATION OF SLAGS FROM STEELMAKING INDUSTRY FOR
PRODUCTION OF SYNTHETIC CALCIUM CARBONATE
As the steel industry accounts for approximately 7% of the total anthropogenic CO2
emissions to the atmosphere (Kim and Worrell, 2002), meaningful reductions may be
achieved by using steelmaking slags for carbon dioxide mineralisation. Due to a high
calcium content, steelmaking slags appear very suitable for this method. The global
theoretical CO2 emissions reduction potential of steelmaking slag carbonation is only
170 Mt/year (Eloneva, 2008) but reductions can be significant for individual steel mills.
The end product of steelmaking slag carbonation could be quite pure calcium carbonate,
if calcium would be extracted from the slags prior to carbonation. Calcium carbonate
has wide markets and special precipitated calcium carbonate (PCC) has a price of >
100€/t (Nordkalk Oy, 2004), making the approach presented here an interesting option.
By utilizing steelmaking slags for production of marketable calcium carbonate and
using it to replace some of the natural and synthetic CaCO3 used in industry, savings in
natural resources would be combined with CO2 emissions reduction.
While various different mineral carbonation processes have been proposed, an
economical and environmentally friendly process that could produce pure calcium
carbonate with a low energy penalty has not yet been found. We have previously
assessed the possibility to use an acetic acid process route, proposed by Kakizawa et al.
(2001) for steelmaking slag carbonation (Teir et al., 2007; Eloneva et al., 2008a, b). It
was found that acetic acid is efficient for dissolving steelmaking slags (Teir et al., 2007;
Eloneva et al., 2008b), but precipitation of calcium carbonate from the resulting solution
required addition of NaOH to allow carbonate precipitation (Eloneva et al., 2008a, b).
This made the process too expensive and energy intensive. An effective calciumselective solvent that can be fully recovered and reused seems to be the key issue for
developing a feasible process for producing pure calcium carbonate from steelmaking
slags. Here, development work on producing pure calcium carbonate from steelmaking
slag by carbonation is presented. Selective extraction of calcium from the slag is
experimentally investigated using various solvents, and precipitation of calcium
carbonate from dissolved calcium was also experimentally studied.
3.1
Dissolution of calcium from steelmaking slags
3.1.1
Solvent selection
In order to find a suitable solvent for selective extraction of calcium from the
steelmaking slag a range of relatively common (weak and strong) acids, bases and salts,
as well as few other solvents were tested. Batches of steel converter slag (74 – 125 µm),
provided by Raahe Works, were used for extraction tests; after filtration the
concentrations of Ca, Si, Fe, Mg, Al, Mn, V and Cr in the filtered solutions were
measured with ICP-AES and AAS. The steel converter slag sample used was also
analysed by total digestion, ICP-AES, AAS, XRF and XRD.
The results from the XRF and ICP-AES analyses of steelmaking slag are listed in
Tables 2a and 2b. The sieved fraction was found to quite well represent the batch.
9/15
According to the XRD analysis, the major phases in steel converter slag are calcium
iron oxide (Ca2Fe2O5), larnite (Ca2SiO4), lime (CaO), calcium hydroxide (Ca(OH)2),
quartz (SiO2), and wuestite (FeO) (in decreasing order).
Table 2a. XRF analysis of the steel converter slag batch, 0-1 mm Units in wt-% (±0.1 %).
(elements larger than 1.0 wt-% listed).
Ca
32.8
(±0.7)
Fe
16.6
(±0.2)
Si
6.5
(±0.1)
F
6.1
(±0.5)
Mg
2.22
(±0.06)
Mn
2.3
(±0.1)
Al
1.97
(±0.08)
V
0.98
(±0.06)
Table 2b. ICP-AES analysis of sieved fraction (74-125 µm) of steel converter slag (average
of 4 sets, only selected elements measured. Units in wt-% (±0.1 %).
Ca
30.1
Fe
14.4
Si
5.0
Mg
2.8
Mn
2.2
Al
1.2
V
1.1
Cr
0.2
As shown in Figure 5, calcium was dissolved efficiently (> 50 %) from the steel
converter slag in most of the acids and ammonium salt solutions. 85-100 % of the slag’s
calcium dissolved in all these acids with concentration greater than 0.5 M, while the
strongest salt solutions were able to dissolve ~70-80 % of the calcium. Acetic acid
dissolved calcium most efficiently. Calcium dissolved quite similarly in solutions of
ammonium chloride and ammonium acetate, but ammonium nitrate was clearly the most
efficient salt solution dissolving ~10 %-units more calcium than ammonium chloride or
ammonium acetate.
100
CH3COOH
HNO3
CH3CH2COOH
Percent of calcium extracted (%)
80
NH4NO3
NH4Cl
CH3COONH4
60
Al(NO3)3
(NH4)2SO4
H2SO4
40
NaCl
(NH2)2CO
CH3COONa
Al2(SO4)3
20
NH4H2PO4
(NH4)2HPO4
NaOH
0
0
0.5
1
1.5
2
Solvent concentration (mol/l)
Figure 5. Dissolution of calcium from steel converter slag (1 g) in various concentrations of
different solvents (50 ml) at room temperature.
Besides calcium extraction the selectivity for this element is important as well. While
only weak concentrations of acids seem to be selective for calcium, all concentrations of
10/15
ammonium salt solutions (0.5-2 M) dissolved calcium selectively from the steel
converter slag. Finally, besides the ability to dissolve calcium selectively from the slag,
the solution should also be favourable for precipitation of calcium carbonate i.e. it
should be alkaline before being exposed to CO2. It was found that the solutions that
dissolve calcium selectively from the steel converter slag seem to result in an alkaline
solution (pH > 7). More detail is given elsewhere (Eloneva 2008 a,b,c).
The results given above were obtained at room temperature; testing for the range 30 70°C with aqueous solutions of ammonium acetate and ammonium chloride gave
increased extraction with increased temperature. Particle size had an effect as well, with
smaller particles giving more extracted calcium for a given time.
Besides steel converter slag, also slags referred to as desulphurisation slag and from
ladle slag (provided by Raahe Works) were used in extraction tests with acetic acid and
ammonium salts at 30 °C. The results are shown in Figure 6.
Percent of calcium extracted (%)
100
CH3COOH_L
CH3COOH_D
80
CH3COONH4_L
CH3COONH4_D
60
NH4NO3_L
NH4NO3_D
40
NH4Cl_L
NH4Cl_D
20
0
0.0
0.5
1.0
1.5
2.0
Concentration (mol/l)
Figure 6. Dissolution of calcium from the ladle slag (pointed by the letter “L” after used
solvent) and from the desulphurization slag (pointed by the letter “D” after used solvent)
(1 g) in various concentrations of different solvents (50 ml) at room temperature.
Thus, the ammonium salts ammonium acetate, ammonium nitrate and ammonium
chloride are all three successful in selectively and efficiently extracting calcium from
steel slags. The main role of these salts is to increase the solubility of calcium which, in
water, would be limited by the solubility of calcium hydroxide. Besides the benefit of
selectivity when compared to acetic acid and other acids, the resulting solution does not
need the use of (expensive) NaOH or other pH modification chemicals that immediately
render the processes uneconomical. One of current activities involves the optimisation
of a continuous system – see Figure 7 - with minimum losses of the solvent chemicals
used, i.e. the ammonium salt. Some testing with recycled solvent liquid was done,
showing that losses of NH3 can be of the order of 10-15% per cycle. This implies that
improved recovery of vaporised NH3 is essential or that a make-up stream of NH3
(rather than the ammonium salt itself) is necessary.
11/15
Residue
Steel slag
Ammonium salt
(make-up)
Carbonation
Extraction
CO2
CaCO3
Figure 8. Schematic of a continuous process for CaCO3 production from steel slags.
3.2
Precipitation of calcium carbonate
3.2.1
General
Since formation of acid during carbonation of solution produced from the steelmaking
slag and acid prevents efficient precipitation of calcium carbonate, ammonium salt
solutions seem to be the most promising solvents for dissolving calcium selectively
from the steelmaking slags prior to carbonation. For acetic acid another possibility
could be the use of for example tributyl phosphate (TBP) for extracting the acid from
the water phase as it is produced during carbonation – see Eloneva et al., 2008c.
3.2.2
Carbonation of calcium containing ammonium salt solution
Due to promising results with the ammonium salt solutions and the steel converter slag,
precipitation of calcium carbonate was tested by carbonating aqueous solutions
containing ammonium salts and calcium dissolved from the steel converter slag, at room
temperature (~20°C). First, steel converter slag was extracted with the three ammonium
salts (ammonium acetate, nitrate and chloride). Then the residual slag was removed and
solution samples were taken for Ca analysis (ICP-AES). The filtered solution of
ammonium salt and dissolved steel converter slag was carbonated; the precipitate was
washed, dried and analysed using XRD, XRF, total carbon (TC) and SEM.
The XRD analysis showed that all precipitates consisted of calcium carbonate as calcite.
The XRF and TC analysis confirmed that the major elements of the precipitates were Ca
and C. The sum of other components identified by XRF amounted only to 0.14–0.21 wt%, indicating that the purity of the calcite produced was 99.8 %. The SEM pictures of
the precipitate showed that the precipitates were in the form of rhombohedral calcite
with a diameter of about 5–30 µm – see Figure 9. The total conversion of calcium in the
slag to calcium in the precipitate was 26-32%.
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However, conversions of calcium from calcium containing solutions to calcium in the
precipitate were much higher (~50-70 %), indicating that the calcium extraction had not
been as efficient as with the dissolution experiments. This is probably due to the low
stirring speed during calcium extraction and could be solved by using better equipment
for stirring. Also, as mentioned above, the losses of NH3 from the process lead to
significant amounts of dissolved calcium in the liquid after the carbonation step.
Figure 9. SEM pictures of the precipitates produced from the solution of ammonium salt
and dissolved steel converter slag.
4.
CONCLUSIONS
Mineralisation of CO2 is an interesting (due to its enormous potential at very many
locations) and urgent (due to the slow deployment of geological storage of CO2) option
for CCS that appears to have great potential for countries like Finland where this
appears to be the only CCS option. An overview was given of methods and procedures
for magnesium silicate carbonation, with emphasis on processes that involve a gas/solid
route for the carbonation step, preferably using Mg(OH)2. Based on that, a stepwise
process using a high pressure high temperature fluidised bed is being developed for
carbonation at up to 600°C, 100 bar (allowing for super-critical conditions for CO2).
Besides this, a production route for Mg(OH)2 from serpentinite rock is being optimised.
Also reported here was the production of pure calcium carbonate from steelmaking
slags by carbonation. The key issue is to find a solvent that would dissolve calcium
selectively and efficiently and could be fully recovered and reused. It was found that
acetic acid dissolved calcium most efficiently (~100 %), while ammonium nitrate was
clearly the most efficient salt solution (~ 80 %). The ammonium salt (ammonium
acetate, ammonium nitrate and ammonium chloride ) solutions were found to dissolve
calcium selectively from the steel converter slag and resulted in an alkaline solution,
likely favourable for precipitation of calcium carbonate. It was found that the higher the
solution temperature, the higher was the ratio of dissolved calcium and the smaller the
slag’s grain size distribution, the higher was the calcium extraction. These results are in
a good agreement with results by Kodama et al. (2008), who also studied
experimentally the dissolution of steel converter slag in ammonium chloride solutions
followed by precipitation of calcium carbonate. Precipitation experiments were carried
out with solutions containing ammonium salt and dissolved steel converter slag as well
as with the solution of acetic acid and dissolved steel converter slag. It was found that
calcium carbonate precipitated as rhombohedral calcite from the solution containing
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ammonium salt and dissolved steel converter slag. Calcium conversion from the
solution into the precipitate was 50-70 %, while the purity of the precipitate was 99.8 %.
The next task is to find out which solvent gives the cheapest and most efficient route to
produce pure calcium carbonate from the steelmaking slags.
5.
ACKNOWLEDGEMENTS
The work on magnesium silicate reported here was done under Ekokem Oy support
funding (2000); the Finnish National research programmes ClimTech (1999-2002) and
ClimBus (2004-2008) by the Finnish Funding Agency for Technology and Innovation
Tekes; Nordic Energy Research (2003-2007); KH Renlund Foundation (2007,2008);
and the Academy of Finland research programme Sustainable Energy (2008-2011). The
work on steel slag carbonation was done under Tekes’ research programme ClimBus
(2004-2008), projects Slag2Pcc (2005-2007) and Slag2Pcc+ (2007-2009).
6.
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Nowadays: Fagerlund, J.
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