Investigations for the environmentally friendly

Journal of Cleaner Production 23 (2012) 195e208

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Investigations for the environmentally friendly production of Na2CO3 and HCl from exhaust CO2, NaCl and H2O Martin Forster* Building Technologies Group, Siemens Schweiz AG, CH-6301 Zug, Switzerland

a r t i c l e i n f o

a b s t r a c t

Article history: Received 12 August 2011 Received in revised form 9 October 2011 Accepted 10 October 2011 Available online 17 October 2011

The conventional Solvay ammonia soda process is a net producer of CO2 and produces large quantities of ecologically doubtful side products. Therefore a possible solution for this problem was investigated. Theoretical and experimental data are given which show the feasibility of a modified ammonia soda process which delivers Na2CO3 and HCl by using exhaust CO2, NaCl and H2O. This modified ammonia soda process would not produce the byproduct CaCl2 as in the conventional Solvay ammonia soda process, would be completely recyclable and could be driven by solar thermal energy. Low maximum reaction temperatures of T 800 K and an estimated achievable solar efficiency of 10% show that this cycle is not only environmentally friendly but also energetically interesting. Kinetic constants of the main reactions are given which are similar to the ones in the conventional process. The principle of a simple solar thermo-chemical reactor is described. Preliminary economical considerations show that this new process might even be competitive when driven by solar thermal energy instead of using fossil fuels. If this novel process would be implemented worldwide approximately up to 3 107 tonne of CO2 could be omitted annually compared with the conventional Solvay ammonia soda process. This would correspond to 0.15% of the annual release of all anthropogenically produced CO2. Ó 2011 Elsevier Ltd. All rights reserved.

Keywords: Modified ammonia soda process CO2 omission Green chemistry Conversion of exhaust CO2 Solar thermo-chemical reaction Environmentally friendly production of Na2CO3

1. Introduction Carbon dioxide (CO2) has been recognized as the main cause of global warming and worldwide efforts to reduce the emission of CO2 and the capture and safe disposal of CO2 are increasingly investigated (Lackner, 2010; Mikkelsen et al., 2010; Zheng et al., 2010). One of the main sources of anthropogenic CO2 is the gaseous exhaust of power plants. By increasing the efficiency of power plants the amount of CO2 (kW h)1 can be reduced (Hong et al., 2009) and by capturing the CO2 produced in power plants (Gauer and Heschel, 2006; Zhang and Lior, 2008; MacDowell et al., 2010; Rivera-Tinoco and Bouallou, 2010; Sayari et al., 2011) and subsequently stored as carbonates the CO2 can be withdrawn from the atmosphere completely (Lackner et al., 1997; Zevenhoven et al., 2008; Eloneva et al., 2008). Another way to reduce the concentration of CO2 in the atmosphere is the conversion of CO2 generally (Schwärzler and Abbreviations: n, number of reaction or equation; DrHT1eT2 (n þ m), sum of DrH0 s of reactions n and m in the temperature interval T1 to T2; DrGx/y(n), DrG of reaction n going from state x to state y; hbasic (n), solar efficiency of reaction n calculated from DrG- and DrH-values; hreactor, efficiency of solar reactor; ton,

tonne ¼ 1000 kg. * Present address: CH-8645 Rapperswil-Jona. Tel.: þ41 55 210 38 44; fax: þ41 41 723 50 32. E-mail address: [email protected]. 0959-6526/$ e see front matter Ó 2011 Elsevier Ltd. All rights reserved. doi:10.1016/j.jclepro.2011.10.012

Schmölz, 1997) or directly from the flue gases of power plants into useful chemical substances by reacting CO2 with a reducing substance at high temperatures and/or with the help of catalysts (Halmann and Steinfeld, 2009) or by electrochemical reduction (Kaneco et al., 2007; Peterson et al., 2010). Photocatalytic and photoelectrochemical reduction of CO2 with water to hydrocarbons on semiconductor materials has been investigated (Adachi et al., 1994; Zhao et al., 2007, 2009) but showed only small efficiencies. The world production of soda (Na2CO3) amounted in 2008 to approximately 4.6 107 ton y1 (Kostick, 2009), whereby 1.2 107 ton y1 of soda were mainly produced from trona (Na3H [CO3]2∙2H2O) and 3.4 107 ton y1 of soda were synthetically produced by the Solvay ammonia soda process. The Solvay ammonia soda process uses NaCl and CaCO3 þ thermal energy and yields Na2CO3 and CaCl2 as byproduct, see Eqs. (1)e(6). Eq. (6) is the net chemical reaction of reactions (1)e(5). In aqueous solution the equilibrium of reaction (6) lies completely on the left hand side and therefore a considerable amount of energy is necessary to drive reactions (1)e(5).

2NH3 þ 2CO2 þ 2H2 O/2NH4 HCO3

(1)

2NH4 HCO3 þ 2NaCl/2NaHCO3 þ 2NH4 Cl

(2)

196

M. Forster / Journal of Cleaner Production 23 (2012) 195e208

2NaHCO3 /Na2 CO3 þ H2 O þ CO2

(3)

CaCO3 /CaO þ CO2

(4)

2NH4 Cl þ CaO/2NH3 þ CaCl2 þ H2 O

(5)

2NaCl þ CaCO3 /Na2 CO3 þ CaCl2

(6)

Reactions (1) and (2) occur at T ¼ 298 . 313 K, reaction (3) at T ¼ 473 K and since reaction (4) needs temperatures up to T ¼ 1323 K and therefore a high amount of thermal energy, which is produced at least partially from carbon based fuel, the Solvay ammonia soda process is a net producer of CO2 besides the wanted product Na2CO3. Furthermore CaCl2 is produced as a byproduct with low economic value, is considered as an environmentally deleterious waste (Kostick, 2009) and is often discarded (Hou, 1942a). Attempts to reduce the negative effects of the byproducts of these reactions have been undertaken (Kasikowski et al., 2004; Gao et al., 2007; Trypuc and Bialowicz, 2011). It seemed therefore interesting to look for a somehow modified ammonia soda process for the production of Na2CO3 without any emission of CO2 and which would omit the occurrence of CaCl2 as a byproduct completely. Furthermore solar thermal energy should be used to drive such a modified ammonia soda process thereby reducing the use of nonrenewable energy sources. Theoretical investigations about possible chemical reactions for such a modified ammonia soda process will be given. All thermodynamic calculations have been performed using literature data (Landolt-Börnstein, 2000) with the simplifications that simultaneously occurring reactions can be treated individually and ideal behavior of the occurring components can be assumed. The most promising chemical reactions have then been investigated experimentally and the data will be discussed. From these information combined with literature data the topology and the positive ecological impact of such a possibly solar driven modified ammonia soda process will be given.

And finally reactions (1)e(3), (7)e(14) could now be replaced by the net chemical reaction (15)

2NaCl þ H2 O þ CO2 /Na2 CO3 þ 2HCl

(15)

As for Eq. (6) also in Eq. (15) the equilibrium lies completely on the left hand side: the reverse reaction of reaction (15) describes the dissolution of soda by hydrogen chloride. Therefore Eq. (15) would also need a considerable amount of energy. But for the modified ammonia soda process this energy should come now from the sun. Me0 or Me00 in Eqs. (7)e(10) and (11)e(14) have now to be chosen in such a way that a) reactions (7)e(14) are thermodynamically favorable and proceed at a reasonable low temperature which can be reached conveniently by a solar thermal reactor. Therefore reactions (7)e(14) should proceed at temperatures T 1273 K and should have DrG 60 kJ at this temperature in order to proceed with a reasonable reaction rate. Similar conditions have been found sufficient for other thermo solar reactions (Forster, 2004). b) Me0 or Me00 should form an oxide which yields an aqueous solution with pH > 7 in order to drive reactions (11)e(14) to the right hand side. c) for this modified ammonia soda process becoming economically favorable Me0 or Me00 have to be chosen such that the corresponding chloride would be cheap and would be a stable material. d) the corresponding chloride should not be poisonous. The only metals which probably can fulfill conditions a)ed) are the alkali metals Li, Na, K and the earth alkali metals Mg and Ca (Srsalts are too expensive and soluble Ba-salts are poisonous). Therefore the thermodynamics of Eqs. (7)e(14) for these different metals were investigated theoretically.

2. Theory 2.1. Principle of a modified ammonia soda process

240

In order to modify reactions (1)e(5) in such a way that no CO2 has to be produced from calcining CaCO3 and no CaCl2 is produced as a byproduct, Eqs. (4) and (5) have to be replaced. Since one mole of CO2 will be necessary for the production of 1 mol Na2CO3 also in a modified ammonia soda process, this CO2 will be taken from the flue gas of a power plant. Therefore Eq. (4) will be obsolete in such a modified ammonia soda process. Reactions (4) and (5) would now have to be replaced by one of the reactions (7)e(10) and one of the reactions (11)e(14), where Me0 and Me00 denote metals which can form mono- and divalent metal cations, respectively:

2Me0 Cl þ H2 O/Me02 O þ 2HCl

(7)

2Me0 Cl þ 2H2 O/2Me0 OH þ 2HCl

(8)

Me00 Cl2 þ H2 O/MeO þ 2HCl

(9)

00

220 200 180 160 140 120 100 80 60

Me Cl2 þ 2H2 O/MeðOHÞ2 þ2HCl

(10)

2NH4 Cl þ Me02 O/2NH3 þ 2MeCl þ H2 O

(11)

2NH4 Cl þ 2Me0 OH/2NH3 þ 2MeCl þ 2H2 O

(12)

2NH4 Cl þ Me00 O/2NH3 þ MeCl2 þ H2 O

(13)

2NH4 Cl þ Me00 ðOHÞ2 /2NH3 þ MeCl2 þ 2H2 O

(14)

40 20 0 300 400 500 600 700 800 900 1000 110012001300 14001500 1600 1700

Temperature / K

Fig. 1. Thermodynamics of Eqs. (6) and (15).


2.2. Thermodynamic data of a modified ammonia soda process Fig. 1 shows DrG as a function of temperature T for reactions (6) and (15). Fig. 1 clearly shows that both reactions can not proceed since DrG >> 0 for 298 K T 1700 K. Fig. 2 shows DrG as a function of temperature T for several reactions (7)e(10) and for different metals Me0 and Me00 and different pressures. Similar calculations for KCl were not performed since the hydrolysis of KCl with H2O vapor occurs at even higher temperatures than with NaCl (Briner and Roth, 1948). According to Fig. 2 the only reaction which can proceed with DrG 20 kJ at T < 1100 K and p ¼ 1 bar is Eq. (9) with Me00 ¼ Mg, which will be denoted as Eq. (9a):

MgCl2 þ H2 O/MgO þ 2HCl

(9a)

In order to get a modified ammonia soda process which is completely recyclable also the corresponding Eq. (13a) with Me00 ¼ Mg

2NH4 Cl þ MgO/2NH3 þ MgCl2 þ H2 O

(13a)

197

has to be thermodynamically favorable and should proceed at a reasonable low temperature. Fig. 3 shows DrG as a function of temperature T for Eqs. (9a) and (13a). Obviously reaction (9a) has DrG 20 kJ already at T ¼ 1000 K and has an equilibrium constant K > 1 for T >¼ 850 K, what makes this reaction suitable for a solar thermal plant. However reaction (9a) proceeds in two steps (Neumann et al., 1935), with Eq. (16) at 623 K (Neumann et al., 1935) or even at 573 K (Gray et al., 2008) and Eq. (17) beginning at about 649 K (Kashani-Nejad et al., 2005):

MgCl2 þ H2 O/MgðOHÞCl þ HCl

(16)

MgðOHÞCl/MgO þ HCl

(17)

According to Fig. 3 reaction (13a) could proceed at the even lower temperature T ¼ 630 K with DrG 20 kJ. Also this temperature would be within easy reach for a high temperature solar reactor. By using pebbles containing MgO, KCl, CaCO3 and caoline which were heated to T > 620 K the vapor of NH4Cl could be decomposed

Fig. 2. Thermodynamics of Eqs. (7)e(10).

198


into NH3 (Solvay, 1886). Further heating these pebbles, now containing MgCl2 according to Eq. (13a), to T ¼ 820 K and introducing water vapor and/or an inert gas like CO2, the pebbles released HCl according to Eq. (9a) (Solvay, 1886, 1891). On the other hand if heated to too high temperatures MgO becomes insoluble in water (Holleman and Wiberg, 1971a) and then Eq. (13a) could not occur at all. Surprisingly reaction (13a) seems to occur already at T 373 K in aqueous solution (Ainscow and Gadgil, 1988), but on the other hand reaction (13a) in aqueous solution would be impossibly slow for practical applications due to a 5∙105 times smaller solubility product of Mg(OH)2 compared to Ca(OH)2 at þ18 C (Hou, 1942b). Because of these different contradictory data it seemed appropriate to investigate reactions (9a) and (13a) experimentally to control at which temperatures these reactions can form a closed cycle (18)

MgCl2 ./MgO./MgCl2 ./ MgO.

(18)

If the reactions of cycle (18) can be shown to proceed with a high yield, and since reactions (1)e(3) are already industrially established for the normal Solvay ammonia soda process, this would prove reaction (15) to be possible. Reaction (15) would now be the overall reaction of the following MgCl2/MgO-modified ammonia soda process:

2NH3 þ 2CO2 þ 2H2 O/2NH4 HCO3

(1)

2NH4 HCO3 þ 2NaCl/2NaHCO3 þ 2NH4 Cl

(2)

2NaHCO3 /Na2 CO3 þ H2 O þ CO2 MgCl2 þ H2 O/MgO þ 2HCl 2NH4 Cl þ MgO/2NH3 þ MgCl2 þ H2 O 2NaCl þ H2 O þ CO2 /Na2 CO3 þ 2HCl

(15)

Analytical grade MgCl2∙6H2O (coarse) and MgO (Sigmae Aldrich), analytical grade NH4Cl, CaO, indicator bromothymol blue and 1N HCl (Omikron) and analytical grade 7N NaOH (Hänseler) were used as received. Coarse MgCl2∙6H2O was used since in real applications also coarse materials would occur. Weighing occurred with a balance (Mettler-Toledo) calibrated to 0.0002 g. Thermo-chemical experiments of Eq. (9a) were performed with an oven 1 (Heraeus), equipped with a quartz tube 2 of 34 mm inner diameter and 1000 mm length, see Fig. 4. Both ends of the quartz tube were equipped with flanges 3 from stainless steel with inner Teflon lining. The flanges could be heated to T > 373 K in order to

100 80 60 40 20

J

(13a)

3. Experimental setup

120

0 -20 -40 -60 -80 -100 -120 -140 350

(9a)

with CO2 taken from the flue gases of a power plant and NaCl taken from the effluent of a desalination plant using sea water. Since all these reactions should proceed at a maximum temperature of T ¼ 1000 K they all could be driven by solar thermal energy collected with a concentrating system.

140

-160 300

(3)

400 450 500 550 600 650 700

750 800

Temperature / K

0 kJ Fig. 3. Thermodynamics of Eqs. (9a) and (13a).

850 900 950 1000


199

Fig. 4. Experimental setup for the reaction of MgCl2∙6H2O with H2O.

avoid condensation. A stainless steel tube 4 with small holes 5 at the end reached into the hot zone of the oven and delivered water vapor by pumping liquid distilled water 6 into the stainless steel tube by an automatic syringe. The flow of water could be regulated from 12 to 120 mL h1, yielding a velocity of water vapor within the quartz tube from 1 to 20 cm s1 depending on the temperature from 473 K to 1073 K and on the flow rate. A non glazed porcelain boat 7 with inner dimensions W H L ¼ 7 5 70 mm and containing MgCl2∙6H2O (1 g) was placed in the quartz tube in the center of the oven. Experiments were done with MgCl2∙6H2O loose or slightly compacted in the porcelain boat or with pieces (8 4 4 mm3) of MgCl2∙6H2O with crystalline density of r ¼ 1.57 g cm3, obtained by pressing powdered MgCl2∙6H2O with F ¼ 40 kN cm2 during 10 min. For reactions with larger amounts of MgCl2∙H2O (12e18 g), two crucibles of aluminum oxide with inner dimensions W H L ¼ 18 18 44 mm were used with MgCl2∙6H2O either loose or slightly compacted. The small holes 5 of the stainless steel tube 4 looked up- and backwards delivering a smooth flow of water vapor 8 above the sample holder 7 toward the exit flange. At the exit flange two cold traps 9 in series condensed the vapors from the thermo-chemical reaction and in an adjoined washing flask 10 with distilled water eventually gaseous HCl was collected. Temperatures at the point of the sample and of both flanges were measured with stainless steel thermocouples of type K, at the point of the sample the thermocouple was surrounded by a ceramic tube of Degussit. Temperature data were collected digitally. Water vapor was produced in the quartz tube at temperatures T > 423 K, temperature rise and fall at the position of the sample were þ18 K min1 and 18 K min1, respectively and the reaction temperature could be stabilized to 3 K. The condensate (HCl) of a thermo-chemical reaction according to Eq. (9a) was titrated with 1N NaOH to pH ¼ 7 (bromothymol blue) and the MgO formed was weighed. From these data the yield of the reaction was calculated. For kinetic measurements the cold traps 9 and the washing flask 10 were replaced by a cooler, as sample holder a thin platinum foil

and samples of MgCl2∙6H2O (1 g) were used. For an experiment the platinum foil was held outside the oven but inside the quartz tube 2 until reaction temperature and water vapor flow had been stabilized. Then the sample was shifted quickly into the middle of the oven. As measured with a thin thermocouple the platinum foil reached the reaction temperature within 25 s thereby the starting point of reaction (9a) could be determined. To the condensates aliquots of 1N NaOH (1 mL) were added as soon as pH ¼ 7 (bromothymol blue) had been reached. From the time of these additions the kinetics of Eq. (9a) could be determined. Reactions of NH4Cl (4 g) with MgO or CaO (some excess compared to theory (Hou, 1942c)), see Eqs. (5) and (13a), were performed in distilled water (25 mL), which was heated to boiling temperature and through which a stream of air (57 mL min1) was pumped, transporting the evolved NH3 into a washing flask filled with water and 1N HCl (10 mL). As soon as pH ¼ 7 (bromothymol blue) had been reached another aliquot of 1N HCl (10 mL) was added to the washing flask etc. until no NH3 was evolved anymore. From the time of these additions the kinetics of Eq. (5) or (13a) could be determined. Grain size distributions were determined with calibrated sieves. Confidential intervals of experimental data have been calculated from experimental uncertainties or from statistical levels of confidence. 4. Experimental results and discussion 4.1. Hydrolysis of MgCl2∙6H2O In run 1 the minimum necessary time to drive Eq. (9a) to completeness using loose MgCl2∙6H2O was determined by varying the reaction time t at T ¼ 1073 K from 32 down to 1 min with 30 mL h1 H2O, see Fig. 5. The error bars represent the estimated experimental errors and of the time during which the oven was held at the specific temperature. Fig. 5 shows that the yield hHCl for the production of HCl with respect to the starting material

200


Fig. 5. Yield hHCl from reaction (9a) as f(t) at T ¼ 1073 K, loose MgCl2∙6H2O; different pretreatments of MgCl2∙6H2O.

MgCl2∙6H2O according to Eq. (9a) has a value of hHCl >¼ 97% for reaction times t ¼ 1e32 min and is constant within error limits. Data at t ¼ 4 min show the repeatability of the experiments. Therefore a reaction time t ¼ 4 min at T ¼ 1073 K is sufficient to drive Eq. (9a) nearly quantitatively to the right hand side. Experiments with loose MgCl2∙6H2O produced some loss of MgO, whereas slightly compacted MgCl2∙6H2O gave no loss. hHCl of reactions with t ¼ 4 min at T ¼ 1073 K with loose, slightly compacted and dense MgCl2∙6H2O were equal within error limits, see Fig. 5. Therefore in further experiments, where the amount of MgO formed was important, MgCl2∙6H2O was slightly compacted. In run 2 the reaction time t ¼ 4 min was held constant and the reaction temperature was changed successively from T ¼ 1073 K down to 473 K with 30 mL h1 H2O, see Fig. 6. Fig. 6 shows that already at T ¼ 800 K a yield hHCl > 97% is achieved for Eq. (9a). By going down to T ¼ 600 K the yield hHCl goes down to approx. 50% corresponding to the production of Mg(OH)Cl according to Eq. (16). Below T ¼ 500 K the yield hHCl goes down to 0% at about T ¼ 450 K. Obviously for such experimental conditions reaction (9a) proceeds nearly quantitatively already at T ¼ 773 K. Fig. 6 shows also the relative weights of the reaction products with respect to the starting material MgCl2∙6H2O. For MgO 19.8% and for Mg(OH)Cl 37.8% would be expected. Obviously MgO was formed down until 770 K and between 750 K and 600 K Mg(OH)Cl was produced. Below 600 K reaction (16) proceeds only partially. All these findings correspond with the formation of HCl. In run 3 MgCl2∙6H2O (12e18 g total) in two crucibles of aluminum oxide were used for each experiment and the temperature was kept to T ¼ 798 K. Table 1 shows the experimental conditions and results. In run 3 for the larger samples of MgCl2∙6H2O with a reaction temperature T ¼ 798 K, t ¼ 7.5 min and 30 mL h1 H2O only a yield hHCl ¼ 72% was achieved, if the starting material MgCl2∙6H2O was slightly compacted in the aluminum oxide crucibles, see Table 1.

With the same experimental conditions, but with loose MgCl2∙6H2O, the yield rose to hHCl ¼ 81.1%. For longer reaction times up to 30 min, loose starting material and a high water flow a nearly quantitative yield hHCl ¼ 98.1% was obtained. Generally enhancing the flow of H2O also enhanced hHCl. The high concentration of 5.7 mol L1 HCl for t ¼ 30 min and 12 mL h1 H2O corresponds to a relation of 8.7 mol H2O per mol of HCl. Fig. 7 shows the grain size distributions of the reaction products of Table 1 together with reactants. From the reaction of slightly compacted reactant only large lumps were obtained and no grain size distribution was possible. From reactions of loose reactant and low hHCl the products still slightly aggregate possibly due to non reacted magnesium chloride hydrates or Mg(OH)Cl and therefore larger particles dominate, whereas in reactions with high hHCl pure MgO is formed yielding smaller particles with the same grain size distribution as commercial analytical grade MgO. In run 4 the kinetics of the formation of HCl from Eq. (9a) for loose MgCl2∙6H2O on a platinum foil with 60 mL h1 H2O was investigated as a function of time t and temperature T, see Fig. 8. For the sake of clearness only 4 data sets out of 8 are shown. Statistical analysis showed that these data could best be fitted with a kinetic first order reaction described by aHCl (t;T) ¼ a(T)(1eb(T)t) with aHCl(t;T) ¼ proportion of HCl developed compared to the theoretical amount possible of reaction (9a) at time t and temperature T and with the kinetic constant k(T) ¼ b(T). All fits had a coefficient of determination R2 >¼ 0.986. Fitting the data with contracting disc or contraction sphere equations (Judd and Norris, 1973) gave much lower R2 and therefore such kinetics were not considered further. Although for Eq. (9a) second order kinetics would be expected, for the experiments of run 1e4 with a high excess of H2O Eq. (9a) has to be rewritten as

MgCl2 $6H2 O þ nH2 O/MgO þ 2HCl þ n þ 5ÞH2 O

with n[1

(9b)


201

Fig. 6. Yield hHCl and weight after reaction (9a) as f(T), reaction time t ¼ 4 min, MgCl2∙6H2O slightly compacted.

Therefore the concentration of H2O does not change very much during reaction (9b). Furthermore with a velocity of v ¼ 213 cm s1 (depending on T and mL h1 H2O) water vapor removes HCl quickly from the reaction site. These conditions then reduce the kinetics of Eq. (9b) to a pseudo first order kinetics what already has been verified by the statistical analysis. Fig. 8 clearly shows that after t ¼ 1800 s at T >¼ 786 K reaction Eq. (9b) has produced aHCl(t;T) >¼ 95% of the theoretical amount of HCl. From Fig. 8 the Arrhenius activation energy EA in the vicinity of T ¼ 773 K could be determined and amounted to EA ¼ 47 5 kJmol1, see Fig. 9. The statistical uncertainties for the frequency factor FF were too large and therefore no reasonable FF can be given. EA found here is lower than EA ¼ 65 kJmol1 from the thermal decomposition of Mg(OH)Cl, measured in a similar temperature region and where also first order kinetics had been found (Kashani-Nejad et al., 2005). Mg(OH)Cl is an intermediate in reaction Eq. (9b), see Eqs. 16 and 17. But Eq. (9b) describes the overall reaction starting from MgCl2∙6H2O which will undergo several dehydration and decomposition steps until MgO and HCl are formed. Obviously some of these intermediate steps have rather low activation energies thereby lowering the overall activation energy EA measured here.

The large differences in hHCl between run 1, 3 and 4 show that the experimental conditions have a large influence on the outcome of reaction (9b). This means that Eq. (9b) proceeds far off equilibrium conditions in the temperature range used for run 1,3 and 4 and might be hindered by the transport of gaseous H2O and HCl within and through reactants and products of Eq. (9b). This is also shown by Fig. 2: for a chemical reaction to proceed nearly quantitatively a DrG of 60 kJ is normally assumed (Forster, 2004). However at T ¼ 800 K the basic reaction (9a) has DrG ¼ þ10 kJ. Therefore in order to drive Eq. (9b) at T ¼ 800 K to the right hand side HCl has to be removed from the reaction site as quick as possible. Obviously this can be achieved with a high amount of and a high flow velocity of water vapor and with loose starting material, see Table 1. Therefore reaction (9b) can proceed nearly quantitatively at T ¼ approx. 800 K if the following precautions are taken: high flow velocity of water vapor, loose starting material MgCl2∙6H2O in a layer of less than 20 mm and a reaction time of up to 30 min. These 30 min compare favorably with reaction times of several hours for burning CaCO3 in the conventional Solvay ammonia soda process (Hou, 1942d). Table 1 shows also that quite highly concentrated HCl could be produced with the experiments of run 3. This would be an economic advantage for the production and selling of HCl.

Table 1 Yield hHCl and concentration of HCl from run 3 of reaction (9a) at T ¼ 798 K with up to 18 g MgCl2∙6H2O. H2O [mL h1]

12 30 60

t ¼ 7.5 min, MgCl2∙6H2O slightly compacted

t ¼ 7.5 min, MgCl2∙6H2O loose

t ¼ 15 min, MgCl2∙6H2O loose

t ¼ 30 min, MgCl2∙6H2O loose

hHCl [%]

HCl [mol L1]

hHCl [%]

HCl [mol L1]

hHCl [%]

HCl [mol L1]

hHCl [%]

HCl [mol L1]

72.0

3.5

71.6 81.1 88.2

5.8 3.6 2.1

93.1

3.8

94.0 97.3 98.1

5.7 3.1 1.5

202


Fig. 7. Grain size distributions: MgO from Table 1 with 1) 30 min, 30 mL h1 H2O; 2) 15 min, 30 mL h1 H2O; 3) 7.5 min, 30 mL h1 H2O; 4) 7.5 min, 12 mL h1 H2O; 5) 7.5 min, 60 mL h1 H2O; 6), 7), 8) commercial analytical grade MgO, CaO, MgCl2∙6H2O respectively.

4.2. Reactions of MgO or CaO with NH4Cl From experiments of reaction (13a) the maximum yield hNH3 of NH3 formation was calculated in comparison with the amount of NH4Cl used, see Table 2. Reactants were commercial analytical grade MgO and CaO as well as MgO produced in reaction (9b) from run 4. Fig. 10 shows the proportion aNH3 ðtÞ of NH3 that had evolved at time t in comparison with the amount of NH4Cl originally present for reaction (13a) or (5) from the experiments of Table 2. The data have then been fitted to the function aNH3 ðtÞ ¼ að1 ebt Þ with the kinetic constant k ¼ b. Here again a pseudo first order kinetics describes reaction (13a): MgO reacts with H2O to Mg(OH)2 which has a limited solubility product and therefore delivers a constant

Fig. 8. Kinetic data of the proportion aHCl (t;T) of HCl evolved from reaction (9b) with 60 mL h1 H2O as f(T,t), loose MgCl2∙6H2O; R2 ¼ coefficient of determination.

concentration of OH during the reaction. OH reacts with NHþ 4 to NH3 and H2O, and NH3 is constantly eliminated by the stream of air. Table 2 clearly shows that the maximum yields hNH3 in reactions (5) and (13a) at T ¼ 371 2 K are equal within error limits for all experiments. Fig. 10 and Table 2 with the kinetics of this reaction show that commercial analytical grade MgO develops NH3 much slower than commercial analytical grade CaO. However two samples of MgO produced from reaction (9b) deliver NH3 at least as fast as commercial analytical grade CaO. Interestingly MgO produced from reaction (9b) with reaction conditions to give the highest hHCl seems also to be the most reactive in Eq. (13a). The reason for this high activity is not yet understood and is the subject of further investigations.

Fig. 9. Determination of the Arrhenius activation energy EA for Eq. (9b) from the data of Fig. 8.


203

Table 2 Yields hNH3 of reactions (13a) or (5) at T ¼ 371 2 K; 1e3 and 7e8: commercial analytical grade; 4: run 4, t ¼ 7.5 min, 12 mL h1 H2O; 5: run 4, t ¼ 7.5 min, 30 mL h1 H2O; 6: run 4, t ¼ 15 min, 30 mL h1 H2O; yields hHCl of reaction (9b); kinetic constants from Fig. 9. No.

1

2

3

4

5

6

7

8

Metaloxide hNH3 [%] hHCl [%] k [s1]

MgO 96.4

MgO 95.6

MgO 97.0

MgO 95.6 71.6 3.9E-04 4E-05

MgO 97.8 81.1 7.2E-04 7E-05

MgO 98.0 93.1 1.4E-03 1E-04

CaO 96.7

CaO 97.5

3.0E-04 2E-05

As a whole the results obtained here confirm the data of Ainscow and Gadgil (1988) and show that the information from Hou (1942b) is definitely outdated: MgO is able to replace CaO in deliberating NH3 from NH4Cl in aqueous solution at 371 2 K with respect to the maximum yield hNH3 and even with respect to kinetics if, at least, the MgO is prepared as given in Chapter 4.1. 4.3. Cycle 18 The foregoing results show that reactions (9b) and (13a) proceed with rather high yields already using a simple equipment. For a sophisticated industrial chemical plant yields near to 100% will then be in easy reach. Furthermore the reaction rate of Eq. (13a) with MgO instead of CaO of Eq. (5) is on a similar level if MgO has been prepared by reaction (9b) at a temperature of T ¼ approx. 798 K. Therefore cycle (18) is definitely a closed cycle and so this MgCl2/MgO-modified ammonia soda process could really be an environmentally friendly way to produce Na2CO3 and HCl from NaCl, H2O and CO2. 5. Discussion with respect to the use of solar thermal energy 5.1. Reactions 3, (9b) and (13a) driven by solar thermal energy For the MgCl2/MgO-modified ammonia soda process, reactions (1)e(3) would still be the same as in the original Solvay ammonia

Fig. 10. Kinetic data of the proportion aNH3 ðtÞ from the reaction: MeO þ 2NH4Cl / MeCl2 þ 2NH3 þ H2O, T ¼ 371 2 K.

8.5E-04 8E-05

soda process. Reactions (3), (9b) and (13a) would now be driven by concentrated solar thermal energy. Solar thermal reactors to perform chemical reactions in the gas phase at very high temperatures are described (Ozalp et al., 2010). Fig. 11 shows now a possible solar thermal reactor for combined solid state/gas phase reactions for continuous operation at intermediate temperatures which would fulfill the restrictions given in Chapter 4.1 for reaction (9b): focused solar radiation is heating up a tube which rotates slowly and is inclined up to 45 against horizontal. The inner wall of the tube would have to be covered with a layer resistant to the chemicals occurring in reaction (9b). The MgCl2∙6H2O enters at the upper end and water vapor at the lower end of the tube and reaction (9b) takes place within the tube. Because of the slow rotation and the slight inclination of the tube MgCl2∙6H2O is transported through the hot zone of the tube within approx. 30 min. During this time reaction (9b) can proceed to completeness and HCl together with excess water vapor leaves at the upper end and MgO at the lower end of the tube. Temperatures up to 823 K are foreseen for parabolic solar trough collectors (Ungeheuer, 2010). Therefore the necessary T ¼ 798 K for reaction (9b) might well be reached with conventional products in the near future, adapted to the specific needs for the chem. compounds occurring in reaction (9b). With the same solar driven chemical reactor from Fig. 11 also reaction (3) could be driven. For reaction (13a) a conventional chemical reactor heated up by solar thermal energy via a heat exchanger could be used. In order to make best use of the collected solar energy the sensible and latent heat content of the reaction products of reactions (3), (9b) and (13a) would have to be used to heat up the corresponding starting materials. In order to reach the required temperatures for reaction (3) and (9b) a minimum solar power would have to be delivered onto the rotating tube by the solar concentrating system. As solar concentrator a field of heliostats or a parabolic trough concentrator can be envisaged. To reach temperatures of approx. 800 K a solar concentration ratio of up to 100 would be necessary using a selective absorber (Pitz-Paal, 2007). If a field of heliostats would be used to concentrate the solar radiation, varying solar irradiation could be compensated for by using more or less heliostats thereby maintaining a constant solar power on the rotating tube. Depending on the average annual solar irradiation at different locations on earth the necessary amount of heliostats will vary. Locations near the equator would be favorites, of course. As with all solar driven thermal processes, for the night time some thermal buffering system has to be provided. Usually molten salt is used for this purpose (Solar Millenium, 2008) which might be used here as well: reaction (3) and (9b), which need higher temperatures, would be driven during day time, additional solar thermal energy would be stored in molten salt to drive reaction (13a) during night time, needing lower temperature. But of course by using solar thermal energy some intermittency of the industrial processes is not avoidable (Nandi and De, 2007; Koroneos et al., 2007; Schnitzer et al., 2007).

204


HCl + H2O MgCl2*6H2O

2° – 45°

H2O MgO

horizontal Fig. 11. Schematic representation of a solar driven reactor to hydrolyze MgCl2∙6H2O.

Fig. 12 gives a pictorial view of the MgCl2/MgO-modified ammonia soda production process.

23H2 OðgÞ/23H2 OðlÞ

5.2. Estimations of the solar efficiency, reduced CO2 emission and economics of such a MgCl2/MgO-modified ammonia soda process

In reaction (9b) a concentration of HCl ¼ 5.55 mol L1 can easily be obtained, see Table 1. This concentration corresponds to a relation of 9 mol of H2O per 1 mol of HCl. Therefore for the real process 1 mol of MgCl2∙6H2O(s) from Eq. (13b) will be used together with 13 mol of H2O for reaction (9b), which can be written as in Eq. (9c)

5.2.1. Solar efficiency Reactions (3), (9a) and (13a) describe basic reactions. In reality reaction (13a) occurs in solution and reaction (9a) delivers a solution and therefore heats of solvation, crystallization etc. will influence the efficiency of this MgCl2/MgO-modified ammonia soda process. From reaction (2) an aqueous solution of 1 mol of NH4Cl in 14 mol of H2O is produced (Hou, 1942e) what is described by Eq. (13b) and subsequently water vapor from Eq. (13b) will be condensed as shown by Eq. (19).

2NH4 Cl ðin 14H2 OÞðlÞ þ MgOðsÞ þ 28H2 OðlÞ/2NH3 ðgÞ þ MgCl2 $6H2 OðsÞ þ 23H2 OðgÞ

(13b)

(19)

MgCl2 $6H2 OðsÞ þ 13H2 OðlÞ/ MgOðsÞ þ 2HCl$9H2 OðlÞ

(9c)

Also reaction (15) is a basic reaction and in reality will deliver aqueous HCl from Eq. (9c) and reads as shown in Eq. (15a):

2NaClðaqÞ þ 19H2 OðlÞ þ CO2 ðgÞ/Na2 CO3 ðsÞ þ 2HCl$9H2 OðlÞ (15a) Using literature data (Cerquetti et al., 1968; Holleman and Wiberg, 1971b; Jahn and Wolf, 1993; Landolt-Börnstein, 1976, 2000; Wagman, 1982;), for Eq. (3), (9c), (13b) the necessary enthalpies DrHT1T2 in different temperature regions and for Eq. (15a) the Gibbs free enthalpy DrG298 was calculated with the assumption that latent and sensible heat can be recovered by 75%, see Table 3. As a result in order to produce 1 mol of Na2CO3 an enthalpy DrH273e798 ¼ 887.5 kJ is necessary which has now to be delivered by solar thermal energy. The basic solar efficiency hbasic of reaction (15a) can be calculated (Kräupel and Steinfeld, 2001; Forster, 2004) as in Eq. (20)

hbasic ð15aÞ ¼ Dr Gp/r ð15aÞ=ðDr H298798 ð3 þ 9c þ 13cÞÞ ¼ 0:16

ð20Þ

with r and p denoting reactants and products, respectively, of Eq. (15a) at 298 K. Solar parabolic trough collectors from 1989 showed an optical to thermal efficiency of 80% at T ¼ 663 K (Solarpaces, 1997) what would correspond to hreactor ¼ 0.8 for reactions (3), (9c) and (13b) (Forster, 2004). With the assumption that in the near future a hreactor ¼ 0.75 at T ¼ 798 K might be achieved, the real solar efficiency of reaction (15a) would become

hreal ¼ hbasic hreactor ¼ 0:12 Fig. 12. Pictorial description of the MgCl2/MgO-modified ammonia soda process.

(21)

Since all necessary energy for pumping liquids and gases and for transporting solids is neglected for these calculations, the


205

Table 3 DrH of reactions (3), (4), (5), (9c) and (13b) in different temperature regions with the assumption that 75% of latent and sensible heat can be recovered (also from Eq. (19)) and DrG298 of reaction (15a). Reaction

DrH298373 [kJ]

DrH373473 [kJ]

(3) (9c) (13b) (3) þ (9c) þ (13b) (15a) (4) (5) CaCl2(aq) discarded (5) CaCl2(anhydrous) separated (3) þ (4) þ (5), CaCl2(aq) discarded (3) þ (4) þ (5), CaCl2(anhydrous) separated

21.9 289.4 307.5 575.0

133.5

DrH373798 [kJ]

DrH298798 total [kJ]

DrH2981323 [kJ]

DrH2981323 total [kJ]

DrG298 [kJ]

178.9 133.5

178.9

887.5 137.7 203.8

22.0 373.9 0.1

133.5

203.8

337.3

351.9

133.5

203.8

689.2

practically achievable solar efficiency for reactions (3) þ (9c) þ (13b) is assumed to reach 10%. This value of 10% is independent of varying solar irradiation: hbasic is a value purely defined by thermodynamics and for T ¼ constant, maintained by the solar concentrating system, hreactor will also be constant.

using the actual fossil fuel price. Reactions (1) and (2) are the same for both processes and need not to be considered and all infrastructure of the soda factory will be the same except the solar thermal reactor(s). For this comparison the following information is helpful:

5.2.2. Reduced CO2 emission For the conventional Solvay ammonia soda process similar calculations were performed, see Table 3. Table 3 distinguishes between the case where the byproduct CaCl2(aq) is discarded as a solution and where CaCl2 (anhydrous) is separated and not discarded. Assuming that the necessary enthalpies will be produced by burning heavy oil with 0.28 kg (kW h)1 CO2 reactions (3) þ (5) are the worldwide annual source of 0.3 107 ton y1 of CO2 or 1.0 107 ton y1 of CO2 for CaCl2 discarded or not discarded, respectively. Since the conventional Solvay ammonia soda process needs for one mol of Na2CO3 also one mol of CaO and the production of 1 ton of CaO releases 0.31 ton of CO2 due to the combustion of fossil fuel for heating CaCO3 in Eq. (4) (Halmann and Steinfeld, 2004) reaction (4) in the conventional Solvay ammonia soda process is another source of 0.6 107 ton of CO2 released annually worldwide. Together with reactions (3) and (5) the conventional Solvay ammonia soda process releases annually worldwide 0.9e1.5 107 ton of CO2 depending if CaCl2 is discarded or not. By contrast the MgCl2/MgO-modified ammonia soda process is a CO2 consumer and needs one mol of CO2 for one mol of Na2CO3 produced. If this CO2 would be taken from the exhausts of power plants a further 1.4 107 ton of CO2 could be eliminated worldwide annually. Therefore by switching over to this novel MgCl2/MgOmodified ammonia soda process worldwide one could omit the release of 2.3e2.9 107 ton of CO2 annually. This saving of CO2 would then correspond to approx. 0.12e0.15% of the annual release of all anthropogenically produced CO2 of 2 1010 ton y1 of CO2 (Steinfeld and Thompson, 1994).

- For locations with high solar irradiation industrial process heat at T ¼ 363 K becomes already cheaper when using solar thermal energy instead of thermal energy from fossil fuels (Kahsay et al., 2011). - A thorough financial analysis of the solar thermal production of lime at Treaction ¼ 1300e1600 K (Eq. (4)) has shown that by using solar thermal technology from 2003 the production costs of lime amounted to 128e157 $/ton CaO (Meier et al., 2005). In 2009 lime produced with fossil fuels was sold for already 105 $/ton CaO (Goonan and Miller, 2010).

5.2.3. Economical considerations of the MgCl2/MgO-modified ammonia soda process Table 4 compares the pros and cons of the MgCl2/MgO-modified with the Solvay ammonia soda process. In order to evaluate the economic differences between the conventional Solvay and the MgCl2/MgO-modified ammonia soda process driven by solar thermal energy it is sufficient to compare the energy costs of reactions (3), (4) and (5) from the conventional with the energy costs of reactions (3), (9c) and (13b) from the modified process. Using the reaction enthalpies DrHxy in Table 3 the energy costs of the corresponding reactions were calculated

Table 4 Pros and cons of the MgCl2/MgO-modified ammonia soda process compared with the conventional Solvay ammonia soda process. Properties of the process Solvay ammonia soda process

MgCl2/MgO-modified ammonia soda process

Thermodynamics possible Basic reactions tested Chemical reactor tested

Yes

Yes

Yes Yes, during 120 years

Yes Yes, during 120 years

Yes Yes (without rotation and sun) Yes (relative to Solvay) Yes No

Huge, during 120 years

Negligible

1323 K

798 K

2e5 h Yes, if CaCl2(aq) discarded No, if CaCl2 (anhydrous) separated No, if CaCl2(aq) discarded Yes, if CaCl2 (anhydrous) separated Emits 0.9e1.5 107 ton CO2/y No

0.5 h No

Kinetics of reactions known Reaction efficiency ok Industrial process developed Cost of research, development and production technology so far Maximum reaction temperature Tmax Time at Tmax Ecologically doubtful byproducts

Economically valuable byproducts

CO2 emission worldwide CO2-certificates for omitted CO2?

Yes

Yes, HCl

Reduces 2.3e2.9 107 ton CO2/y Yes

206


- As generally known the costs of solar thermal energy are declining whereas the costs of process heat produced from fossil fuels are rising. Therefore solar thermal energy at T ¼ 800 K for Eq. (9c) will soon become competitive compared to process heat produced from fossil fuels. Furthermore Table 3 shows that for reactions (3), (9c) and (13b) the largest amount of process heat will be necessary in the temperature range of Treaction ¼ 298e373 K and for this part solar thermal energy is already competitive (Kahsay et al., 2011). A rough estimate of the economics of the MgCl2/MgO-modified ammonia soda process and compared with the conventional Solvay ammonia soda process can therefore be given in Fig. 13 with the following assumptions and data: - The solar installations for the soda factory will be erected at the time when the conventional installations need refurbishments anyway. Infrastructure costs and lifetimes for these solar installations are assumed to be comparable to conventional ones therefore no infrastructure costs need to be considered. - The actual price of Na2CO3 of 213 Euro/ton (Kostick, 2011) produced by the Solvay process is assumed to reflect the true costs for the production of Na2CO3 without subsidies from selling of byproducts, i.e. this means CaCl2(aq) discarded. - To be on the safe side solar thermal energy costs in the temperature ranges 298e373 K, 373e473 K and 373e798 K were assumed to be higher by 20%, 40% and 60%, respectively, than the actual value for fossil fuel of 0.038 Euro/kW h (for manufacturing industries without Climate Change Levy) (Quarterly Energy Prices, 2011). For reactions (3)e(5) energy costs from fossil fuels of 0.038 Euro/kW h were used. For all thermal processes an efficiency of 0.8 was assumed (Meier et al., 2005).

- The following costs and reductions were neglected: concentrating HCl (approx. 18.6%) to HCl (33%), other expenses specifically necessary for the use of solar thermal energy, reductions due to scale; purifying the distiller waste from the Solvay ammonia soda process to obtain pure CaCl2, pollution fees for discarding CaCl2(aq). - As financial data were used: 100 Euro/ton HCl (33%) (ICIS Pricing, 2011), 165 Euro/ton CaCl2 (94e97%, anhydrous, min. 15 ton) (Alibaba, 2011), CO2-certificates of 12 Euro/ton CO2 (European Energy Exchange, 2011). Fig. 13 clearly shows: - The energy costs of the Solvay ammonia soda process constitute 33% of the overall costs of the Na2CO3-production (Trypuc and Bialowicz, 2011). In Fig. 13 for “Solvay, CaCl2(aq) discarded” “a” contains only the thermal energy costs of reactions (3) þ (4) þ (5) which amount to 20% of the overall costs (¼“c”). If the energy costs for reactions (1) and (2) and all other energy costs for driving the soda factory would be added to “a” one would easily end up with the 33% mentioned by Trypuc and Bialowicz (2011). This shows that the calculations for Fig. 13 have a sound basis. - The costs of Na2CO3 when produced with this MgCl2/MgOmodified ammonia soda process driven by solar thermal energy and including selling the byproduct HCl are already lower than the costs of Na2CO3 produced by the conventional Solvay ammonia soda process with CaCl2(aq) discarded. This would correspond to a worldwide saving of 3.8 109 Euro/y. - If all CaCl2(aq) produced in the conventional Solvay process is transformed into CaCl2 (94e97%, anhydrous) and is also included in the comparison, the solar process using rather conservative assumptions is slightly (20%) more expensive. This would correspond to worldwide additional costs of

Fig. 13. Comparison of estimated production costs of Na2CO3 for different scenarios with: a ¼ energy costs for Eq. (3) þ (4) þ (5) with CaCl2(aq) discarded, b ¼ costs for all other reactions þ fix costs etc., c ¼ a þ b ¼ actual costs of Na2CO3 produced by the Solvay ammonia soda process with CaCl2(aq) discarded, d ¼ energy costs for Eq. (3) þ (4) þ (5) with CaCl2 (94e97%, anhydrous) separated, e ¼ income from CaCl2 (94e97%, anhydrous) sold, f ¼ d þ b þ e ¼ overall costs for Na2CO3 including CaCl2 (94e97%, anhydrous) sold, g ¼ solar thermal energy costs for Eq. (3) þ (9c) þ (13b), h ¼ income from HCl (33%) sold, i ¼ income from CO2-certificates, j ¼ g þ b þ h þ i ¼ overall costs for Na2CO3 produced by the MgCl2/ MgO-modified ammonia soda process using solar thermal energy including CO2-certificates and HCl (33%) sold; for assumptions and neglects see text.


5.9 108 Euro/y. But since never 100% of CaCl2(aq) from the Solvay process can be converted into pure CaCl2 the solar process seems to be nearly competitive even for this case. As a whole this new MgCl2/MgO-modified ammonia soda process driven by solar thermal energy seems also to fulfill the economic conditions to be implemented in reality. However if the Soda production industry will really change to this ecologically friendly MgCl2/MgO-modified ammonia soda process depends also on the necessary research and investigations into this new technology. And investigations for new and benign processes have often difficulties to become realized as can be seen with the cement industry (Moya et al., 2011). 6. Conclusion Thermodynamic calculations revealed a possible MgCl2/MgOmodified ammonia soda process and experiments showed that such a novel process can indeed be envisaged. The new process uses the completely closed cycle MgCl2./MgO./MgCl2./MgO. instead of the nonrecycable reactions of burning CaCO3 and discarding the unwanted CaCl2 as with the traditional Solvay ammonia soda process. By contrast this MgCl2/MgO-modified ammonia soda process is environmentally friendly and delivers HCl as a more valuable byproduct. Additionally the kinetics of these new reactions is at least as fast as the conventional one. Furthermore it was shown that temperatures of only 800 K are sufficient to drive this closed cycle. Such temperatures are easily obtained with concentrated solar energy. A solar thermo-chemical reactor was proposed to perform the thermal hydrolysis of MgCl2∙6H2O and the thermal decomposition of NaHCO3 driven by concentrated solar thermal energy. If the sensible and latent heat of all occurring reactions could be used to 75% the practically achievable solar efficiency was estimated to be on the order of 10%. Economic estimates indicated the possibility to drive this new process by solar thermal energy in a competitive way compared to using fossil fuels. If the existing conventional Solvay ammonia soda process would be replaced worldwide by this MgCl2/MgO-modified ammonia soda process and if the necessary CO2 would be taken from the exhausts of power plants the emission of approx. 2.3e2.9 107 ton of CO2 could be omitted annually. This saving of CO2 would then correspond to 0.12e0.15% of the annual release of all anthropogenically produced CO2. Acknowledgements Support from and helpful discussions with Dr. Georges Tenchio, Siemens Schweiz AG, are gratefully acknowledged. This work was supported by Siemens Schweiz AG. References Adachi, K., Ohta, K., Mizuno, T., 1994. Photocatalytic reduction of carbon dioxide to hydrocarbon using copper-loaded titanium dioxide. Solar Energy 53, 187e190. Ainscow, W.S., Gadgil, B.B., 1988, Process for producing magnesium oxide, US Patent 4720375. Alibaba, 2011, alibaba.com/showroom/calcium-chloride-94-97.html, (last accessed 07.10.11.). Briner, E., Roth, P., 1948. Investigations of the hydrolysis of alkaline metal chlorides by water vapor alone or with several other components (in French). Helv. Chim. Acta 31, 1352e1360. Cerquetti, A., Longhi, P., Mussini, T., 1968. Thermodynamics of aqueous hydrochloric acid from E.M.F’s of hydrogen-chlorine cells. J. Chem. Eng. Data 13, 458e461. Eloneva, S., Teir, S., Salminen, J., Fogelholm, C.-J., Zevenhoven, R., 2008. Fixation of CO2 by carbonating calcium derived from blast furnace slag. Energy 33, 1461e1467. European Energy Exchange, 2011, www.eex.com/en/Market%20Data/Trading% 20Data/Emission%20Rights/EU%20Emission%20Allowances%23166%3B%20Spot, (last accessed 17.09.11.).

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