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Journal of Cleaner Production 23 (2012) 195e208

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Journal of Cleaner Production

journal homepage: www.elsevier .com/locate/ jc lepro

Investigations for the environmentally friendly production of Na2CO3and 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

Article history:Received 12 August 2011Received in revised form9 October 2011Accepted 10 October 2011Available online 17 October 2011

Keywords:Modified ammonia soda processCO2 omissionGreen chemistryConversion of exhaust CO2

Solar thermo-chemical reactionEnvironmentally friendly production ofNa2CO3

Abbreviations: n, number of reaction or equationDrH0s of reactions n and m in the temperature intervreaction n going from state x to state y; hbasic (n),calculated from DrG- and DrH-values; hreactor, effictonne ¼ 1000 kg.* Present address: CH-8645 Rapperswil-Jona. Tel.: þ

723 50 32.E-mail address: martin.forster@gmx.ch.

0959-6526/$ e see front matter � 2011 Elsevier Ltd.doi:10.1016/j.jclepro.2011.10.012

a b s t r a c t

The conventional Solvay ammonia soda process is a net producer of CO2 and produces large quantities ofecologically 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 sodaprocess which delivers Na2CO3 and HCl by using exhaust CO2, NaCl and H2O. This modified ammoniasoda process would not produce the byproduct CaCl2 as in the conventional Solvay ammonia sodaprocess, would be completely recyclable and could be driven by solar thermal energy. Low maximumreaction temperatures of T � 800 K and an estimated achievable solar efficiency of 10% show that thiscycle is not only environmentally friendly but also energetically interesting. Kinetic constants of the mainreactions are given which are similar to the ones in the conventional process. The principle of a simplesolar thermo-chemical reactor is described. Preliminary economical considerations show that this newprocess might even be competitive when driven by solar thermal energy instead of using fossil fuels. Ifthis novel process would be implemented worldwide approximately up to 3 � 107 tonne of CO2 could beomitted annually compared with the conventional Solvay ammonia soda process. This would correspondto 0.15% of the annual release of all anthropogenically produced CO2.

� 2011 Elsevier Ltd. All rights reserved.

1. Introduction

Carbon dioxide (CO2) has been recognized as the main cause ofglobalwarming andworldwide efforts to reduce the emission of CO2and the capture and safe disposal of CO2 are increasingly investi-gated (Lackner, 2010;Mikkelsen et al., 2010; Zheng et al., 2010). Oneof the main sources of anthropogenic CO2 is the gaseous exhaust ofpower plants. By increasing the efficiency of power plants theamount of CO2 (kW h)�1 can be reduced (Hong et al., 2009) and bycapturing the CO2 produced in power plants (Gauer and Heschel,2006; Zhang and Lior, 2008; MacDowell et al., 2010; Rivera-Tinocoand Bouallou, 2010; Sayari et al., 2011) and subsequently stored ascarbonates the CO2 can be withdrawn from the atmospherecompletely (Lackner et al., 1997; Zevenhoven et al., 2008; Elonevaet al., 2008). Another way to reduce the concentration of CO2 inthe atmosphere is the conversion of CO2 generally (Schwärzler and

; DrHT1eT2 (n þ m), sum ofal T1 to T2; DrGx/y(n), DrG ofsolar efficiency of reaction niency of solar reactor; ton,

41 55 210 38 44; fax: þ41 41

All rights reserved.

Schmölz, 1997) or directly from the flue gases of power plants intouseful chemical substances by reacting CO2 with a reducingsubstance 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 pho-toelectrochemical reduction of CO2 with water to hydrocarbons onsemiconductor 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 toapproximately 4.6 � 107 ton y�1 (Kostick, 2009), whereby1.2 � 107 ton y�1 of soda were mainly produced from trona (Na3H[CO3]2∙2H2O) and 3.4 � 107 ton y�1 of soda were syntheticallyproduced by the Solvayammonia soda process. The Solvayammoniasoda process uses NaCl and CaCO3 þ thermal energy and yieldsNa2CO3 and CaCl2 as byproduct, see Eqs. (1)e(6). Eq. (6) is the netchemical reaction of reactions (1)e(5). In aqueous solution theequilibrium of reaction (6) lies completely on the left hand side andtherefore a considerable amount of energy is necessary to drivereactions (1)e(5).

2NH3 þ 2CO2 þ 2H2O/2NH4HCO3 (1)

2NH4HCO3 þ 2NaCl/2NaHCO3 þ 2NH4Cl (2)

240

180

200

220

140

160

180

100

120

60

80

0

20

40

Temperature / K

300 400 500 600 700 800 900 1 000 1100 1200 1 300 1 400 1 500 1 600 1 700

0

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

M. Forster / Journal of Cleaner Production 23 (2012) 195e208196

2NaHCO3/Na2CO3 þ H2Oþ CO2 (3)

CaCO3/CaOþ CO2 (4)

2NH4Clþ CaO/2NH3 þ CaCl2 þH2O (5)

2NaClþ CaCO3/Na2CO3 þ CaCl2 (6)

Reactions (1) and (2) occur at T ¼ 298 . 313 K, reaction (3) atT ¼ 473 K and since reaction (4) needs temperatures up toT¼ 1323 K and therefore a high amount of thermal energy, which isproduced at least partially from carbon based fuel, the Solvayammonia soda process is a net producer of CO2 besides the wantedproduct Na2CO3. Furthermore CaCl2 is produced as a byproductwith low economic value, is considered as an environmentallydeleterious waste (Kostick, 2009) and is often discarded (Hou,1942a). Attempts to reduce the negative effects of the byproductsof 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 modifiedammonia soda process for the production of Na2CO3 without anyemission of CO2 and which would omit the occurrence of CaCl2 asa byproduct completely. Furthermore solar thermal energy shouldbe used to drive such a modified ammonia soda process therebyreducing the use of nonrenewable energy sources.

Theoretical investigations about possible chemical reactions forsuch a modified ammonia soda process will be given. All thermo-dynamic calculations have been performed using literature data(Landolt-Börnstein, 2000) with the simplifications that simulta-neously occurring reactions can be treated individually and idealbehavior of the occurring components can be assumed. The mostpromising chemical reactions have then been investigated experi-mentally and the data will be discussed. From these informationcombined with literature data the topology and the positiveecological impact of such a possibly solar drivenmodified ammoniasoda process will be given.

2. Theory

2.1. Principle of a modified ammonia soda process

In order to modify reactions (1)e(5) in such a way that no CO2has to be produced from calcining CaCO3 and no CaCl2 is producedas a byproduct, Eqs. (4) and (5) have to be replaced. Since one moleof CO2 will be necessary for the production of 1 mol Na2CO3 also ina modified ammonia soda process, this CO2 will be taken from theflue gas of a power plant. Therefore Eq. (4) will be obsolete in sucha modified ammonia soda process.

Reactions (4) and (5) would now have to be replaced by one ofthe reactions (7)e(10) and one of the reactions (11)e(14), whereMe0 and Me00 denote metals which can form mono- and divalentmetal cations, respectively:

2Me0Clþ H2O/Me02Oþ 2HCl (7)

2Me0Clþ 2H2O/2Me0OHþ 2HCl (8)

Me00Cl2 þ H2O/MeOþ 2HCl (9)

Me00Cl2 þ 2H2O/MeðOHÞ2þ2HCl (10)

2NH4ClþMe02O/2NH3 þ 2MeClþH2O (11)

2NH4Clþ 2Me0OH/2NH3 þ 2MeClþ 2H2O (12)

2NH4ClþMe00O/2NH3 þMeCl2 þH2O (13)

2NH4ClþMe00ðOHÞ2/2NH3 þMeCl2 þ 2H2O (14)

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

2NaClþ H2Oþ CO2/Na2CO3 þ 2HCl (15)

As for Eq. (6) also in Eq. (15) the equilibrium lies completely onthe left hand side: the reverse reaction of reaction (15) describesthe dissolution of soda by hydrogen chloride. Therefore Eq. (15)would also need a considerable amount of energy. But for themodified ammonia soda process this energy should come now fromthe sun.

Me0 or Me00 in Eqs. (7)e(10) and (11)e(14) have now to bechosen in such a way that

a) reactions (7)e(14) are thermodynamically favorable andproceed at a reasonable low temperature which can be reachedconveniently by a solar thermal reactor. Therefore reactions(7)e(14) should proceed at temperatures T � 1273 K andshould have DrG � �60 kJ at this temperature in order toproceed with a reasonable reaction rate. Similar conditionshave been found sufficient for other thermo solar reactions(Forster, 2004).

b) Me0 or Me00 should form an oxide which yields an aqueoussolutionwith pH> 7 in order to drive reactions (11)e(14) to theright hand side.

c) for this modified ammonia soda process becoming economi-cally favorable Me0 or Me00 have to be chosen such that thecorresponding chloride would be cheap and would be a stablematerial.

d) the corresponding chloride should not be poisonous.

The only metals which probably can fulfill conditions a)ed) arethe alkali metals Li, Na, K and the earth alkali metals Mg and Ca (Sr-salts are too expensive and soluble Ba-salts are poisonous).Therefore the thermodynamics of Eqs. (7)e(14) for these differentmetals were investigated theoretically.

M. Forster / Journal of Cleaner Production 23 (2012) 195e208 197

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 proceedsince DrG >> 0 for 298 K � T � 1700 K.

Fig. 2 shows DrG as a function of temperature T for severalreactions (7)e(10) and for different metals Me0 and Me00 anddifferent pressures. Similar calculations for KCl were not performedsince the hydrolysis of KCl with H2O vapor occurs at even highertemperatures than with NaCl (Briner and Roth, 1948).

According to Fig. 2 the only reaction which can proceed withDrG � �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 H2O/MgOþ 2HCl (9a)

In order to get a modified ammonia soda process which iscompletely recyclable also the corresponding Eq. (13a) withMe00 ¼ Mg

2NH4ClþMgO/2NH3 þMgCl2 þH2O (13a)

Fig. 2. Thermodynamic

has to be thermodynamically favorable and should proceed ata 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 atT 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 at573 K (Gray et al., 2008) and Eq. (17) beginning at about 649 K(Kashani-Nejad et al., 2005):

MgCl2 þH2O/MgðOHÞClþ HCl (16)

MgðOHÞCl/MgOþHCl (17)

According to Fig. 3 reaction (13a) could proceed at the evenlower temperature T ¼ 630 K with DrG � �20 kJ. Also thistemperature would be within easy reach for a high temperaturesolar reactor.

By using pebbles containing MgO, KCl, CaCO3 and caoline whichwere heated to T > 620 K the vapor of NH4Cl could be decomposed

s of Eqs. (7)e(10).

M. Forster / Journal of Cleaner Production 23 (2012) 195e208198

into NH3 (Solvay, 1886). Further heating these pebbles, now con-taining MgCl2 according to Eq. (13a), to T ¼ 820 K and introducingwater vapor and/or an inert gas like CO2, the pebbles released HClaccording to Eq. (9a) (Solvay, 1886, 1891).

On the other hand if heated to too high temperatures MgObecomes insoluble in water (Holleman and Wiberg, 1971a) andthen Eq. (13a) could not occur at all. Surprisingly reaction (13a)seems to occur already at T � 373 K in aqueous solution (Ainscowand Gadgil, 1988), but on the other hand reaction (13a) inaqueous solution would be impossibly slow for practical applica-tions due to a 5∙105 times smaller solubility product of Mg(OH)2compared to Ca(OH)2 at þ18 �C (Hou, 1942b).

Because of these different contradictory data it seemed appro-priate to investigate reactions (9a) and (13a) experimentally tocontrol at which temperatures these reactions can form a closedcycle (18)

MgCl2./MgO./MgCl2./MgO. (18)

If the reactions of cycle (18) can be shown to proceedwith a highyield, and since reactions (1)e(3) are already industrially estab-lished for the normal Solvay ammonia soda process, this wouldprove reaction (15) to be possible.

Reaction (15) would now be the overall reaction of the followingMgCl2/MgO-modified ammonia soda process:

2NH3 þ 2CO2 þ 2H2O/2NH4HCO3 (1)

2NH4HCO3 þ 2NaCl/2NaHCO3 þ 2NH4Cl (2)

120

140

60

80

100

120

J

20

40

60

J

-4 0

-2 0

0

-1 00

-8 0

-6 0

-1 60

-1 40

-1 20

Tempe

300 350 400 450 500 550 600

0 kJ

Fig. 3. Thermodynamics o

2NaHCO3/Na2CO3 þH2Oþ CO2 (3)

MgCl2 þH2O/MgOþ 2HCl (9a)

2NH4ClþMgO/2NH3 þMgCl2 þ H2O (13a)

2NaClþ H2Oþ CO2/Na2CO3 þ 2HCl (15)

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

3. Experimental setup

Analytical grade MgCl2∙6H2O (coarse) and MgO (SigmaeAldrich), analytical grade NH4Cl, CaO, indicator bromothymol blueand 1N HCl (Omikron) and analytical grade 7N NaOH (Hänseler)were used as received. Coarse MgCl2∙6H2O was used since in realapplications also coarse materials would occur. Weighing occurredwith a balance (Mettler-Toledo) calibrated to �0.0002 g.

Thermo-chemical experiments of Eq. (9a) were performed withan oven 1 (Heraeus), equipped with a quartz tube 2 of 34 mm innerdiameter and 1000 mm length, see Fig. 4. Both ends of the quartztube were equipped with flanges 3 from stainless steel with innerTeflon lining. The flanges could be heated to T > 373 K in order to

rature / K

650 700 750 800 850 900 950 1000

f Eqs. (9a) and (13a).

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

M. Forster / Journal of Cleaner Production 23 (2012) 195e208 199

avoid condensation. A stainless steel tube 4 with small holes 5 atthe end reached into the hot zone of the oven and delivered watervapor by pumping liquid distilled water 6 into the stainless steeltube by an automatic syringe. The flow of water could be regulatedfrom 12 to 120 mL h�1, yielding a velocity of water vapor within thequartz tube from 1 to 20 cm s�1 depending on the temperaturefrom 473 K to 1073 K and on the flow rate.

A non glazed porcelain boat 7 with inner dimensionsW � H � L ¼ 7 � 5 � 70 mm and containing MgCl2∙6H2O (1 g) wasplaced in thequartz tube in the centerof theoven. ExperimentsweredonewithMgCl2∙6H2O loose or slightly compacted in the porcelainboat or with pieces (8� 4� 4mm3) of MgCl2∙6H2Owith crystallinedensity of r ¼ 1.57 g cm�3, obtained by pressing powderedMgCl2∙6H2Owith F¼ 40 kN cm�2 during 10min. For reactions withlarger amounts ofMgCl2∙H2O (12e18 g), two crucibles of aluminumoxide with inner dimensions W � H � L ¼ 18 � 18 � 44 mm wereusedwithMgCl2∙6H2O either loose or slightly compacted. The smallholes 5 of the stainless steel tube 4 looked up- and backwardsdelivering a smoothflowofwater vapor 8 above the sample holder 7toward the exit flange. At the exit flange two cold traps 9 in seriescondensed the vapors from the thermo-chemical reaction and in anadjoined washing flask 10 with distilled water eventually gaseousHCl was collected. Temperatures at the point of the sample and ofboth flanges were measured with stainless steel thermocouples oftype K, at the point of the sample the thermocouplewas surroundedby a ceramic tube of Degussit. Temperature data were collecteddigitally. Water vapor was produced in the quartz tube at temper-atures T > 423 K, temperature rise and fall at the position of thesample were þ18 K min�1 and �18 K min�1, respectively and thereaction temperature could be stabilized to �3 K. The condensate(HCl) of a thermo-chemical reaction according to Eq. (9a) wastitrated with 1N NaOH to pH ¼ 7 (bromothymol blue) and the MgOformed was weighed. From these data the yield of the reaction wascalculated.

For kinetic measurements the cold traps 9 and the washing flask10 were replaced by a cooler, as sample holder a thin platinum foil

and samples of MgCl2∙6H2O (1 g) were used. For an experiment theplatinum foil was held outside the oven but inside the quartz tube 2until reaction temperature and water vapor flow had been stabi-lized. Then the sample was shifted quickly into the middle of theoven. As measured with a thin thermocouple the platinum foilreached the reaction temperature within 25 s thereby the startingpoint of reaction (9a) could be determined. To the condensatesaliquots of 1N NaOH (1 mL) were added as soon as pH ¼ 7 (bro-mothymol blue) had been reached. From the time of these addi-tions the kinetics of Eq. (9a) could be determined.

Reactions of NH4Cl (4 g) with MgO or CaO (some excesscompared to theory (Hou, 1942c)), see Eqs. (5) and (13a), wereperformed in distilled water (25 mL), which was heated to boilingtemperature and through which a stream of air (57 mL min�1) waspumped, transporting the evolved NH3 into a washing flask filledwith water and 1N HCl (10 mL). As soon as pH ¼ 7 (bromothymolblue) had been reached another aliquot of 1N HCl (10 mL) wasadded 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 calculatedfrom experimental uncertainties or from statistical levels ofconfidence.

4. Experimental results and discussion

4.1. Hydrolysis of MgCl2∙6H2O

In run 1 the minimum necessary time to drive Eq. (9a) tocompleteness using loose MgCl2∙6H2O was determined by varyingthe reaction time t at T ¼ 1073 K from 32 down to 1 min with30 mL h�1 H2O, see Fig. 5. The error bars represent the estimatedexperimental errors and of the time during which the oven washeld at the specific temperature. Fig. 5 shows that the yield hHCl forthe production of HCl with respect to the starting material

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

M. Forster / Journal of Cleaner Production 23 (2012) 195e208200

MgCl2∙6H2O according to Eq. (9a) has a value of hHCl >¼ 97% forreaction 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 todrive Eq. (9a) nearly quantitatively to the right hand side.

Experiments with loose MgCl2∙6H2O produced some loss ofMgO, whereas slightly compacted MgCl2∙6H2O gave no loss. hHCl ofreactions with t ¼ 4 min at T ¼ 1073 K with loose, slightly com-pacted and dense MgCl2∙6H2O were equal within error limits, seeFig. 5. Therefore in further experiments, where the amount of MgOformed was important, MgCl2∙6H2O was slightly compacted.

In run 2 the reaction time t ¼ 4 min was held constant and thereaction temperature was changed successively from T ¼ 1073 Kdown to 473 K with 30 mL h�1 H2O, see Fig. 6. Fig. 6 shows thatalready at T ¼ 800 K a yield hHCl > 97% is achieved for Eq. (9a). Bygoing 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) proceedsnearly quantitatively already at T ¼ 773 K.

Fig. 6 shows also the relative weights of the reaction productswith respect to the starting material MgCl2∙6H2O. For MgO 19.8%and for Mg(OH)Cl 37.8% would be expected. Obviously MgO wasformed down until 770 K and between 750 K and 600 K Mg(OH)Clwas 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 ofaluminum oxide were used for each experiment and the temper-ature was kept to T ¼ 798 K. Table 1 shows the experimentalconditions and results.

In run 3 for the larger samples of MgCl2∙6H2O with a reactiontemperature T ¼ 798 K, t ¼ 7.5 min and 30 mL h�1 H2O only a yieldhHCl ¼ 72% was achieved, if the starting material MgCl2∙6H2O wasslightly compacted in the aluminum oxide crucibles, see Table 1.

With the same experimental conditions, but with looseMgCl2∙6H2O, the yield rose to hHCl ¼ 81.1%. For longer reactiontimes up to 30 min, loose starting material and a high water flowa nearly quantitative yield hHCl ¼ 98.1% was obtained. Generallyenhancing the flow of H2O also enhanced hHCl. The high concen-tration of 5.7 mol L�1 HCl for t ¼ 30 min and 12 mL h�1 H2Ocorresponds to a relation of 8.7 mol H2O per mol of HCl.

Fig. 7 shows the grain size distributions of the reaction productsof Table 1 together with reactants. From the reaction of slightlycompacted reactant only large lumps were obtained and no grainsize distribution was possible. From reactions of loose reactant andlow hHCl the products still slightly aggregate possibly due to nonreacted magnesium chloride hydrates or Mg(OH)Cl and thereforelarger particles dominate, whereas in reactions with high hHCl pureMgO is formed yielding smaller particles with the same grain sizedistribution as commercial analytical grade MgO.

In run 4 the kinetics of the formation of HCl from Eq. (9a) forloose MgCl2∙6H2O on a platinum foil with 60 mL h�1 H2O wasinvestigated as a function of time t and temperature T, see Fig. 8. Forthe sake of clearness only 4 data sets out of 8 are shown. Statisticalanalysis showed that these data could best be fitted with a kineticfirst order reaction described by aHCl (t;T) ¼ a(T)(1�e�b(T)t) withaHCl(t;T)¼ proportion of HCl developed compared to the theoreticalamount possible of reaction (9a) at time t and temperature T andwith the kinetic constant k(T) ¼ b(T). All fits had a coefficient ofdetermination R2>¼ 0.986. Fitting the datawith contracting disc orcontraction sphere equations (Judd and Norris, 1973) gave muchlower R2 and therefore such kinetics were not considered further.Although for Eq. (9a) second order kinetics would be expected, forthe experiments of run 1e4with a high excess of H2O Eq. (9a) has tobe rewritten as

MgCl2$6H2Oþ nH2O/MgOþ 2HClþ �

nþ 5ÞH2O with n[1 (9b)

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

M. Forster / Journal of Cleaner Production 23 (2012) 195e208 201

Therefore the concentration of H2O does not change very muchduring reaction (9b). Furthermorewith a velocity of v¼ 2�13 cm s�1

(depending on T and mL h�1 H2O) water vapor removes HCl quicklyfrom the reaction site. These conditions then reduce the kinetics ofEq. (9b) to a pseudo first order kinetics what already has beenverified by the statistical analysis.

Fig. 8 clearly shows that after t ¼ 1800 s at T >¼ 786 K reactionEq. (9b) has produced aHCl(t;T)>¼ 95% of the theoretical amount ofHCl.

From Fig. 8 the Arrhenius activation energy EA in the vicinity ofT ¼ 773 K could be determined and amounted toEA ¼ 47 � 5 kJmol�1, see Fig. 9. The statistical uncertainties for thefrequency factor FF were too large and therefore no reasonable FFcan be given. EA found here is lower than EA ¼ 65 kJmol�1 from thethermal decomposition of Mg(OH)Cl, measured in a similartemperature region and where also first order kinetics had beenfound (Kashani-Nejad et al., 2005). Mg(OH)Cl is an intermediate inreaction Eq. (9b), see Eqs. 16 and 17.

But Eq. (9b) describes the overall reaction starting fromMgCl2∙6H2O which will undergo several dehydration and decom-position steps until MgO and HCl are formed. Obviously some ofthese intermediate steps have rather low activation energiesthereby lowering the overall activation energy EA measured here.

Table 1Yield hHCl and concentration of HCl from run 3 of reaction (9a) at T ¼ 798 K with up to 1

H2O [mL h�1] t ¼ 7.5 min, MgCl2∙6H2Oslightly compacted

t ¼ 7.5 min, MgCl2∙6H2Oloose

hHCl [%] HCl [mol L�1] hHCl [%] HCl [mol

12 71.6 5.830 72.0 3.5 81.1 3.660 88.2 2.1

The large differences in hHCl between run 1, 3 and 4 show that theexperimental conditions have a large influence on the outcome ofreaction (9b). This means that Eq. (9b) proceeds far off equilibriumconditions in the temperature rangeused for run1,3 and4andmightbe hindered by the transport of gaseous H2O and HCl within andthrough reactants and products of Eq. (9b). This is also shown byFig. 2: for a chemical reaction to proceed nearly quantitatively a DrGof � �60 kJ is normally assumed (Forster, 2004). However atT¼ 800K the basic reaction (9a) hasDrG¼þ10 kJ. Therefore in orderto drive Eq. (9b) at T ¼ 800 K to the right hand side HCl has to beremoved from the reaction site as quick as possible. Obviously thiscan be achieved with a high amount of and a high flow velocity ofwater vapor and with loose starting material, see Table 1.

Therefore reaction (9b) can proceed nearly quantitatively atT ¼ approx. 800 K if the following precautions are taken: high flowvelocity of water vapor, loose starting material MgCl2∙6H2O ina layer of less than 20 mm and a reaction time of up to 30 min.These 30 min compare favorably with reaction times of severalhours for burning CaCO3 in the conventional Solvay ammonia sodaprocess (Hou, 1942d).

Table 1 shows also that quite highly concentrated HCl could beproduced with the experiments of run 3. This would be aneconomic advantage for the production and selling of HCl.

8 g MgCl2∙6H2O.

t¼ 15 min, MgCl2∙6H2O loose t¼ 30 min, MgCl2∙6H2O loose

L�1] hHCl [%] HCl [mol L�1] hHCl [%] HCl [mol L�1]

94.0 5.793.1 3.8 97.3 3.1

98.1 1.5

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

M. Forster / Journal of Cleaner Production 23 (2012) 195e208202

4.2. Reactions of MgO or CaO with NH4Cl

From experiments of reaction (13a) the maximum yield hNH3of

NH3 formation was calculated in comparison with the amount ofNH4Cl used, see Table 2. Reactants were commercial analyticalgrade MgO and CaO as well as MgO produced in reaction (9b) fromrun 4.

Fig. 10 shows the proportion aNH3ðtÞ of NH3 that had evolved at

time t in comparison with the amount of NH4Cl originally presentfor reaction (13a) or (5) from the experiments of Table 2. The datahave then been fitted to the function aNH3

ðtÞ ¼ að1� e�btÞwith thekinetic constant k ¼ b. Here again a pseudo first order kineticsdescribes reaction (13a): MgO reacts with H2O to Mg(OH)2 whichhas 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) with60 mL h�1 H2O as f(T,t), loose MgCl2∙6H2O; R2 ¼ coefficient of determination.

concentration of OH� during the reaction. OH� reacts with NHþ4 to

NH3 and H2O, and NH3 is constantly eliminated by the stream of air.Table 2 clearly shows that the maximumyields hNH3

in reactions(5) and (13a) at T ¼ 371 � 2 K are equal within error limits for allexperiments. Fig. 10 and Table 2 with the kinetics of this reactionshow that commercial analytical grade MgO develops NH3 muchslower than commercial analytical grade CaO. However twosamples of MgO produced from reaction (9b) deliver NH3 at least asfast as commercial analytical grade CaO. Interestingly MgOproduced from reaction (9b) with reaction conditions to give thehighest hHCl seems also to be the most reactive in Eq. (13a). Thereason for this high activity is not yet understood and is the subjectof further investigations.

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

Table 2Yields hNH3

of reactions (13a) or (5) at T¼ 371�2 K; 1e3 and 7e8: commercial analytical grade; 4: run 4, t¼ 7.5min,12mL h�1 H2O; 5: run 4, t¼ 7.5min, 30mL h�1 H2O; 6: run4, t ¼ 15 min, 30 mL h�1 H2O; yields hHCl of reaction (9b); kinetic constants from Fig. 9.

No. 1 2 3 4 5 6 7 8

Metaloxide MgO MgO MgO MgO MgO MgO CaO CaOhNH3

[%] 96.4 95.6 97.0 95.6 97.8 98.0 96.7 97.5hHCl [%] 71.6 81.1 93.1k [s�1] 3.0E-04 � 2E-05 3.9E-04 � 4E-05 7.2E-04 � 7E-05 1.4E-03 � 1E-04 8.5E-04 � 8E-05

M. Forster / Journal of Cleaner Production 23 (2012) 195e208 203

As a whole the results obtained here confirm the data ofAinscow and Gadgil (1988) and show that the information fromHou (1942b) is definitely outdated: MgO is able to replace CaO indeliberating NH3 from NH4Cl in aqueous solution at 371 � 2 K withrespect to the maximum yield hNH3

and even with respect tokinetics 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) proceedwith rather high yields already using a simple equipment. Fora sophisticated industrial chemical plant yields near to 100% willthen 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 MgOhas been prepared by reaction (9b) at a temperature of T ¼ approx.798 K. Therefore cycle (18) is definitely a closed cycle and so thisMgCl2/MgO-modified ammonia soda process could really be anenvironmentally friendly way to produce Na2CO3 and HCl fromNaCl, 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 aNH3ðtÞ from the reaction: MeO þ

2NH4Cl / MeCl2 þ 2NH3 þ H2O, T ¼ 371 � 2 K.

soda process. Reactions (3), (9b) and (13a) would now be driven byconcentrated solar thermal energy.

Solar thermal reactors to perform chemical reactions in the gasphase at very high temperatures are described (Ozalp et al., 2010).Fig. 11 shows now a possible solar thermal reactor for combinedsolid state/gas phase reactions for continuous operation at inter-mediate temperatures which would fulfill the restrictions given inChapter 4.1 for reaction (9b): focused solar radiation is heating upa tube which rotates slowly and is inclined up to 45� againsthorizontal. The inner wall of the tube would have to be coveredwith a layer resistant to the chemicals occurring in reaction (9b).The MgCl2∙6H2O enters at the upper end and water vapor at thelower end of the tube and reaction (9b) takes place within the tube.Because of the slow rotation and the slight inclination of the tubeMgCl2∙6H2O is transported through the hot zone of the tube withinapprox. 30 min. During this time reaction (9b) can proceed tocompleteness and HCl together with excess water vapor leaves atthe upper end and MgO at the lower end of the tube.

Temperatures up to 823 K are foreseen for parabolic solar troughcollectors (Ungeheuer, 2010). Therefore the necessary T ¼ 798 K forreaction (9b) might well be reached with conventional products inthe near future, adapted to the specific needs for the chem.compounds occurring in reaction (9b). With the same solar drivenchemical reactor from Fig. 11 also reaction (3) could be driven. Forreaction (13a) a conventional chemical reactor heated up by solarthermal energy via a heat exchanger could be used. In order tomake best use of the collected solar energy the sensible and latentheat content of the reaction products of reactions (3), (9b) and (13a)would have to be used to heat up the corresponding startingmaterials.

In order to reach the required temperatures for reaction (3) and(9b) a minimum solar power would have to be delivered onto therotating tube by the solar concentrating system. As solar concen-trator a field of heliostats or a parabolic trough concentrator can beenvisaged. To reach temperatures of approx. 800 K a solarconcentration ratio of up to 100 would be necessary using a selec-tive absorber (Pitz-Paal, 2007).

If a field of heliostats would be used to concentrate the solarradiation, varying solar irradiation could be compensated for byusing more or less heliostats thereby maintaining a constant solarpower on the rotating tube. Depending on the average annual solarirradiation at different locations on earth the necessary amount ofheliostats will vary. Locations near the equator would be favorites,of course.

As with all solar driven thermal processes, for the night timesome thermal buffering system has to be provided. Usually moltensalt is used for this purpose (SolarMillenium, 2008)whichmight beused here as well: reaction (3) and (9b), which need highertemperatures, would be driven during day time, additional solarthermal energy would be stored in molten salt to drive reaction(13a) during night time, needing lower temperature. But of courseby using solar thermal energy some intermittency of the industrialprocesses is not avoidable (Nandi and De, 2007; Koroneos et al.,2007; Schnitzer et al., 2007).

H + l C H 2 O

l C g M 2

H 6 * 2 O

H2O

° 5 4 – ° 2

O g M

H2 O

l a t n o z i r o h

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

M. Forster / Journal of Cleaner Production 23 (2012) 195e208204

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

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

5.2.1. Solar efficiencyReactions (3), (9a) and (13a) describe basic reactions. In reality

reaction (13a) occurs in solution and reaction (9a) delivers a solu-tion and therefore heats of solvation, crystallization etc. willinfluence the efficiency of this MgCl2/MgO-modified ammonia sodaprocess. From reaction (2) an aqueous solution of 1 mol of NH4Cl in14 mol of H2O is produced (Hou, 1942e) what is described by Eq.(13b) and subsequently water vapor from Eq. (13b) will becondensed as shown by Eq. (19).

2NH4Cl ðin 14H2OÞðlÞ þMgOðsÞ þ 28H2OðlÞ/2NH3ðgÞþMgCl2$6H2OðsÞ þ 23H2OðgÞ (13b)

Fig. 12. Pictorial description of the MgCl2/MgO-modified ammonia soda process.

23H2OðgÞ/23H2OðlÞ (19)

In reaction (9b) a concentration of HCl ¼ 5.55 mol L�1 can easilybe obtained, see Table 1. This concentration corresponds to a rela-tion of 9 mol of H2O per 1 mol of HCl. Therefore for the real process1 mol of MgCl2∙6H2O(s) from Eq. (13b) will be used together with13 mol of H2O for reaction (9b), which can be written as in Eq. (9c)

MgCl2$6H2OðsÞ þ 13H2OðlÞ/MgOðsÞ þ 2HCl$9H2OðlÞ (9c)

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

2NaClðaqÞ þ 19H2OðlÞ þ CO2ðgÞ/Na2CO3ðsÞ þ 2HCl$9H2OðlÞ(15a)

Using literature data (Cerquetti et al., 1968; Holleman andWiberg, 1971b; Jahn and Wolf, 1993; Landolt-Börnstein, 1976,2000; Wagman, 1982;), for Eq. (3), (9c), (13b) the necessaryenthalpies DrHT1�T2 in different temperature regions and for Eq.(15a) the Gibbs free enthalpy DrG298 was calculated with theassumption 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 enthalpyDrH273e798¼ 887.5 kJ is necessary which has now to be delivered bysolar thermal energy.

The basic solar efficiency hbasic of reaction (15a) can be calcu-lated (Kräupel and Steinfeld, 2001; Forster, 2004) as in Eq. (20)

hbasic ð15aÞ ¼ �DrGp/rð15aÞ=ðDrH298�798ð3 þ 9cþ 13cÞÞ¼ 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 tothermal efficiency of 80% at T ¼ 663 K (Solarpaces, 1997) whatwould correspond to hreactor ¼ 0.8 for reactions (3), (9c) and (13b)(Forster, 2004). With the assumption that in the near futurea hreactor ¼ 0.75 at T ¼ 798 K might be achieved, the real solarefficiency of reaction (15a) would become

hreal ¼ hbasic�hreactor ¼ 0:12 (21)

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

Table 3DrH 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)) andDrG298 of reaction (15a).

Reaction DrH298�373 [kJ] DrH373�473 [kJ] DrH373�798 [kJ] DrH298�798 total [kJ] DrH298�1323 [kJ] DrH298�1323 total [kJ] DrG298 [kJ]

(3) �21.9 133.5(9c) 289.4 178.9(13b) 307.5(3) þ (9c) þ (13b) 575.0 133.5 178.9 887.5(15a) 137.7(4) 203.8(5) CaCl2(aq) discarded 22.0(5) CaCl2(anhydrous) separated 373.9(3) þ (4) þ (5), CaCl2(aq)

discarded0.1 133.5 203.8 337.3

(3) þ (4) þ (5), CaCl2(anhydrous)separated

351.9 133.5 203.8 689.2

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

Properties of the process Solvay ammoniasoda process

MgCl2/MgO-modifiedammonia sodaprocess

Thermodynamicspossible

Yes Yes

Basic reactions tested Yes YesChemical reactor tested Yes, during 120 years Yes (without rotation

and sun)Kinetics of reactions

knownYes Yes (relative to

Solvay)Reaction efficiency ok Yes YesIndustrial process

developedYes, during 120 years No

Cost of research,developmentand productiontechnology so far

Huge, during 120 years Negligible

Maximum reactiontemperature Tmax

1323 K 798 K

Time at Tmax 2e5 h 0.5 hEcologically doubtful

byproductsYes, if CaCl2(aq)discarded

No

No, if CaCl2 (anhydrous)separated

Economically valuablebyproducts

No, if CaCl2(aq)discarded

Yes, HCl

Yes, if CaCl2 (anhydrous)separated

CO2 emissionworldwide

Emits 0.9e1.5 � 107 tonCO2/y

Reduces 2.3e2.9 � 107 tonCO2/y

CO2-certificates foromitted CO2?

No Yes

M. Forster / Journal of Cleaner Production 23 (2012) 195e208 205

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

5.2.2. Reduced CO2 emissionFor the conventional Solvay ammonia soda process similar

calculations were performed, see Table 3. Table 3 distinguishesbetween the case where the byproduct CaCl2(aq) is discarded asa solution and where CaCl2 (anhydrous) is separated and not dis-carded. Assuming that the necessary enthalpies will be producedby 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 y�1 of CO2 or1.0 � 107 ton y�1 of CO2 for CaCl2 discarded or not discarded,respectively.

Since the conventional Solvay ammonia soda process needsfor one mol of Na2CO3 also one mol of CaO and the production of1 ton of CaO releases 0.31 ton of CO2 due to the combustion offossil fuel for heating CaCO3 in Eq. (4) (Halmann and Steinfeld,2004) reaction (4) in the conventional Solvay ammonia sodaprocess is another source of 0.6 � 107 ton of CO2 releasedannually worldwide. Together with reactions (3) and (5) theconventional Solvay ammonia soda process releases annuallyworldwide 0.9e1.5 � 107 ton of CO2 depending if CaCl2 is dis-carded or not.

By contrast the MgCl2/MgO-modified ammonia soda process isa CO2 consumer and needs one mol of CO2 for one mol of Na2CO3produced. If this CO2 would be taken from the exhausts of powerplants a further 1.4� 107 ton of CO2 could be eliminated worldwideannually. Therefore by switching over to this novel MgCl2/MgO-modified ammonia soda process worldwide one could omit therelease of 2.3e2.9 � 107 ton of CO2 annually. This saving of CO2would then correspond to approx. 0.12e0.15% of the annual releaseof all anthropogenically produced CO2 of 2 � 1010 ton y�1 of CO2(Steinfeld and Thompson, 1994).

5.2.3. Economical considerations of the MgCl2/MgO-modifiedammonia soda process

Table 4 compares the pros and cons of the MgCl2/MgO-modifiedwith the Solvay ammonia soda process.

In order to evaluate the economic differences between theconventional Solvay and the MgCl2/MgO-modified ammonia sodaprocess driven by solar thermal energy it is sufficient to comparethe energy costs of reactions (3), (4) and (5) from the conventionalwith the energy costs of reactions (3), (9c) and (13b) from themodified process. Using the reaction enthalpies DrHx�y in Table 3the energy costs of the corresponding reactions were calculated

using the actual fossil fuel price. Reactions (1) and (2) are the samefor both processes and need not to be considered and all infra-structure of the soda factory will be the same except the solarthermal reactor(s). For this comparison the following information ishelpful:

- For locations with high solar irradiation industrial process heatat T ¼ 363 K becomes already cheaper when using solarthermal energy instead of thermal energy from fossil fuels(Kahsay et al., 2011).

- A thorough financial analysis of the solar thermal production oflime at Treaction ¼ 1300e1600 K (Eq. (4)) has shown that byusing solar thermal technology from 2003 the production costsof lime amounted to 128e157 $/ton CaO (Meier et al., 2005). In2009 lime produced with fossil fuels was sold for already 105$/ton CaO (Goonan and Miller, 2010).

M. Forster / Journal of Cleaner Production 23 (2012) 195e208206

- As generally known the costs of solar thermal energy aredeclining whereas the costs of process heat produced fromfossil fuels are rising.

Therefore solar thermal energy at T¼ 800 K for Eq. (9c) will soonbecome competitive compared to process heat produced fromfossil fuels. Furthermore Table 3 shows that for reactions (3), (9c)and (13b) the largest amount of process heat will be necessary inthe temperature range of Treaction ¼ 298e373 K and for this partsolar thermal energy is already competitive (Kahsay et al., 2011). Arough estimate of the economics of the MgCl2/MgO-modifiedammonia soda process and compared with the conventional Solvayammonia soda process can therefore be given in Fig. 13 with thefollowing assumptions and data:

- The solar installations for the soda factory will be erected atthe time when the conventional installations need refurbish-ments anyway. Infrastructure costs and lifetimes for thesesolar installations are assumed to be comparable to conven-tional ones therefore no infrastructure costs need to beconsidered.

- The actual price of Na2CO3 of 213 Euro/ton (Kostick, 2011)produced by the Solvay process is assumed to reflect the truecosts for the production of Na2CO3 without subsidies fromselling of byproducts, i.e. this means CaCl2(aq) discarded.

- To be on the safe side solar thermal energy costs in thetemperature ranges 298e373 K, 373e473 K and 373e798 Kwere assumed to be higher by 20%, 40% and 60%, respectively,than the actual value for fossil fuel of 0.038 Euro/kW h (formanufacturing industries without Climate Change Levy)(Quarterly Energy Prices, 2011). For reactions (3)e(5) energycosts from fossil fuels of 0.038 Euro/kW h were used. For allthermal processes an efficiency of 0.8 was assumed (Meieret al., 2005).

Fig. 13. Comparison of estimated production costs of Na2CO3 for different scenarios with: areactions þ fix costs etc., c ¼ a þ b ¼ actual costs of Na2CO3 produced by the Solvay ammonCaCl2 (94e97%, anhydrous) separated, e ¼ income from CaCl2 (94e97%, anhydrous) sold, f ¼ dthermal energy costs for Eq. (3) þ (9c) þ (13b), h ¼ income from HCl (33%) sold, i ¼ income frMgO-modified ammonia soda process using solar thermal energy including CO2-certificate

- The following costs and reductions were neglected: concen-trating HCl (approx. 18.6%) to HCl (33%), other expensesspecifically necessary for the use of solar thermal energy,reductions due to scale; purifying the distiller waste from theSolvay ammonia soda process to obtain pure CaCl2, pollutionfees for discarding CaCl2(aq).

- As financial data were used: 100 Euro/ton HCl (33%) (ICISPricing, 2011), 165 Euro/ton CaCl2 (94e97%, anhydrous, min. 15ton) (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 consti-tute 33% of the overall costs of the Na2CO3-production (Trypucand 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 theenergy costs for reactions (1) and (2) and all other energy costsfor driving the soda factory would be added to “a” one wouldeasily end up with the 33% mentioned by Trypuc and Bialowicz(2011). This shows that the calculations for Fig. 13 have a soundbasis.

- The costs of Na2CO3 when produced with this MgCl2/MgO-modified ammonia soda process driven by solar thermalenergy and including selling the byproduct HCl are alreadylower than the costs of Na2CO3 produced by the conventionalSolvay ammonia soda process with CaCl2(aq) discarded. Thiswould correspond to a worldwide saving of 3.8 � 109 Euro/y.

- If all CaCl2(aq) produced in the conventional Solvay process istransformed into CaCl2 (94e97%, anhydrous) and is alsoincluded in the comparison, the solar process using ratherconservative assumptions is slightly (20%) more expensive.This would correspond to worldwide additional costs of

¼ energy costs for Eq. (3) þ (4) þ (5) with CaCl2(aq) discarded, b ¼ costs for all otheria soda process with CaCl2(aq) discarded, d ¼ energy costs for Eq. (3) þ (4) þ (5) withþ b þ e ¼ overall costs for Na2CO3 including CaCl2 (94e97%, anhydrous) sold, g ¼ solar

om CO2-certificates, j ¼ g þ b þ h þ i ¼ overall costs for Na2CO3 produced by the MgCl2/s and HCl (33%) sold; for assumptions and neglects see text.

M. Forster / Journal of Cleaner Production 23 (2012) 195e208 207

5.9 � 108 Euro/y. But since never 100% of CaCl2(aq) from theSolvay process can be converted into pure CaCl2 the solarprocess seems to be nearly competitive even for this case.

As a whole this new MgCl2/MgO-modified ammonia sodaprocess driven by solar thermal energy seems also to fulfill theeconomic conditions to be implemented in reality. However if theSoda production industry will really change to this ecologicallyfriendly MgCl2/MgO-modified ammonia soda process depends alsoon the necessary research and investigations into this new tech-nology. And investigations for newand benign processes have oftendifficulties to become realized as can be seen with the cementindustry (Moya et al., 2011).

6. Conclusion

Thermodynamic calculations revealed a possible MgCl2/MgO-modified ammonia soda process and experiments showed that sucha novel process can indeed be envisaged. The new process uses thecompletely closed cycle MgCl2./MgO./MgCl2./MgO.instead of the nonrecycable reactions of burning CaCO3 and dis-carding the unwanted CaCl2 as with the traditional Solvay ammoniasoda process. By contrast this MgCl2/MgO-modified ammonia sodaprocess is environmentally friendly and delivers HCl as a morevaluable byproduct. Additionally the kinetics of these new reactionsis at least as fast as the conventional one. Furthermore it was shownthat temperatures of only 800 K are sufficient to drive this closedcycle. Such temperatures areeasilyobtainedwith concentrated solarenergy. A solar thermo-chemical reactor was proposed to performthe thermal hydrolysis of MgCl2∙6H2O and the thermal decompo-sition of NaHCO3 driven by concentrated solar thermal energy. If thesensible and latent heat of all occurring reactions could be used to75% thepracticallyachievable solarefficiencywasestimated tobeonthe order of 10%. Economic estimates indicated the possibility todrive this newprocess by solar thermal energy in a competitivewaycompared to using fossil fuels.

If the existing conventional Solvay ammonia soda process wouldbe replaced worldwide by this MgCl2/MgO-modified ammoniasoda process and if the necessary CO2 would be taken from theexhausts of power plants the emission of approx. 2.3e2.9 � 107 tonof CO2 could be omitted annually. This saving of CO2 would thencorrespond to 0.12e0.15% of the annual release of all anthro-pogenically produced CO2.

Acknowledgements

Support from and helpful discussions with Dr. Georges Tenchio,Siemens Schweiz AG, are gratefully acknowledged. This work wassupported by Siemens Schweiz AG.

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