Chem. Mater. 2010, 22, 851–859 851 DOI:10.1021/cm9016529
Solar Syngas Production via H2O/CO2-Splitting Thermochemical Cycles with Zn/ZnO and FeO/Fe3O4 Redox Reactions† A. Stamatiou,‡ P. G. Loutzenhiser,‡ and A. Steinfeld*,‡,§ ‡
Solar Technology Laboratory, Paul Scherrer Institute, 5232 Villigen PSI, Switzerland, and §Department of Mechanical and Process Engineering, ETH Zurich, 8092 Zurich, Switzerland Received June 15, 2009. Revised Manuscript Received August 9, 2009
The solar production of syngas from H2O and CO2 is examined via two-step thermochemical cycles based on Zn/ZnO and FeO/Fe3O4 redox reactions. The first, endothermic step is the thermal dissociation of the metal oxide using concentrated solar radiation as the energy source of hightemperature process heat. The second, nonsolar, exothermic step is the reaction of the metal or reduced metal oxide with a mixture of H2O and CO2 yielding syngas (H2 and CO), together with the initial form of the metal oxide that is recycled to the first step. Chemical equilibrium compositions for the systems of Zn and FeO with CO2 þ H2O were computed as a function of temperature and pressure for different stoichiometries. A series of dynamic thermogravimetric experimental runs in the range 673-1423 K was carried out to evaluate the reaction kinetics and syngas quality of the second step. The molar flow rate fractions of the gaseous products exhibited linear dependencies on the molar flow rate fractions of the gaseous reactants for both the FeO/Fe3O4 and Zn/ZnO systems. or
1. Introduction Two-step solar thermochemical cycles based on Zn/ZnO and FeO/Fe3O4 redox reactions are proposed for producing synthesis gas (syngas, with H2 and CO as main constituents) from H2O and CO2 as sole chemical sources. The first step is the highly endothermic dissociation of the metal oxide that proceeds at above 2000 K using concentrated solar radiation as the high-temperature heat source, represented as ZnO f Zn þ
1 o O2 ΔH298K ¼ 350:5 kJ mol -1 2
ð1aÞ
or Fe3 O4 f 3FeO þ
1 o ¼ 316:6 kJ mol -1 O2 ΔH298K 2
ð1bÞ
which has been experimentally demonstrated in solar furnaces.1-4 The novel second step of the cycle is the lowtemperature exothermic (nonsolar) reduction of a mixture of CO2 and H2O to syngas, represented as Zn þ βCO2 þ ð1 -βÞH2 O f ZnO þ βCO þ ð1 -βÞH2 (2a) -108:8 kJ mol -1 < ΔH o298K < -67:5 kJ mol -1
† Accepted as part of the 2010 “Materials Chemistry of Energy Conversion Special Issue”. *Author to whom correspondence should be addressed. E-mail: aldo.
[email protected]. ::
(1) Muller, R.; Haeberling, P.; Palumbo, R. D. Sol. Energy 2006, 80, 500–511. (2) Schunk, L; Haeberling, P.; Wepf, S.; Wuillemin, D.; Meier, A.; Steinfeld, A. J. Sol. Energy Eng. 2008, 130, 021009. (3) Schunk, L.; Steinfeld, A. AIChE J. 2009, 55, 1497–1504. (4) Charvin, P.; Abanades, S.; Flamant, G.; Lemort, F. Energy 2007, 32, 1124–1133.
r 2009 American Chemical Society
3FeO þ βCO2 þ ð1 -βÞH2 O f Fe3 O4 þ βCO þ ð1 -βÞH2 (2b) -77:8 kJ mol -1 < ΔH o298K < -33:6 kJ mol -1
with 0 < β < 1. The resulting syngas can be further processed via Fischer-Tropsch to liquid fuels5 or combusted for electricity generation6 in highly efficient combined Brayton-Rankine cycles. The H2/CO and CO2/CO ratios of the syngas are important indicators of its quality for further application. Often, an additional water-gas shift step is added to adjust the desired syngas composition.7 A schematic of the two-step solar-driven thermochemical cycle for a generic metal oxide redox system MxOy/M is depicted in Figure 1. H2O-splitting cycles with metal oxide redox reactions, described in-depth in previous works,8-17 have the potential of reaching high (5) Dry, M. E. Catal. Today 2002, 71, 227–241. (6) Mantzaras, J. Combust. Sci. Technol. 2008, 180, 1137–1168. (7) Wilhelm, D. J.; Simbeck, D. R.; Karp, A. D.; Dickenson, R. L. Fuel Process. Technol. 2001, 71, 139–148. (8) Steinfeld, A. Sol. Energy 2005, 78/5, 603–615. (9) Steinfeld, A. Int. J. Hydrogen Energy 2002, 27, 611–619. (10) Abanades, S.; Charvin, P.; Flamant, G.; Neveu, P. Energy 2006, 31 (14), 2805–2822. (11) Perkins, C.; Weimer, A. W. Int. J. Hydrogen Energy 2004, 29(15), 1587–1599. (12) Perkins, C.; Weimer, A. W. AIChE J. 2009, 55(2), 286–293. (13) Allendorf, M. D.; Diver, R. B.; Siegel, N. P.; Miller, J. E. Energy Fuels 2008, 22, 4115–4124. (14) Kodama, T.; Nakamuro, Y.; Mizuno, T. J. Sol. Energy Eng. 2006, 128(1), 3–7. (15) Abanades, S.; Flamant, G.:: Sol. Energy 2006, 80(12), 1611–1623. (16) Roeb, M.; Sattler, C.; Kluser, R.; Monnerie, N.; Oliveira, L. d.; Konstandopoulos, A. G.; Agrafiotis, C.; Zaspalis, V.; Nalbandian, L.; Steele, A.; Stobbe, P. J. Sol. Energy Eng. 2006, 128(2), 125–133. (17) Abanades, S.; Charvin, P.; Lemont, F.; Flamant, G. Int. J. Hydrogen Energy 2008, 33(21), 6021–6030.
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Figure 1. Schematic of the two-step H2O/CO2-splitting solar thermochemical cycle for syngas production with metal oxide redox reactions.
solar-to-chemical energy conversion efficiencies and, therefore, economic competitiveness vis- a-vis other routes for producing solar fuels. Further, thermochemical cycles allow for the long-term storage and long-range transportation of solar energy, decoupling the intermittency and location of the solar chemical plant from the 24 h-dispatchability of the fuel at the customer site. In Figure 1, the second step of the traditional H2O-splitting cycle is extended to simultaneously split H2O and CO2 and coproduce H2 and CO. In contrast to the direct thermolysis of H2O and CO2, the proposed thermochemical cycle allows for lower maximum temperatures and for the derivation of syngas and O2 in separate steps, thereby bypassing the need for their high-temperature gas separation. The oxidation of Zn using mixtures of H2O/CO2/CO/ H2 as reactive gas was initially studied within the framework of the Imperial Smelting Process.18-21 The increasing interest in H2 production has resulted in new studies using H2O as the oxidant,22,23 and its experimental demonstration for both Zn24-26 and FeO.4 An alternative second step to the hydrolysis is the reduction of captured CO2 to avoid its sequestration and produce carbonneutral solar fuels. The oxidation of Zn, FeO, and other (18) Clarke, J. A.; Fray, D. J. Trans. Inst. Min. Metall. 1979, 88, C161– C166. (19) Lewis, L. A.; Cameron, A. M. Metall. Trans. B 1995, 26, 911–918. (20) Lewis, L. A.; Cameron, A. M. Metall. Trans. B 1995, 26, 919–924. (21) Clarke, J. A.; Fray, D. J. Trans. Inst. Min. Metall. 1982, 91, C26– C31. (22) Vishnevetsky, I.; Epstein, M. Int. J. Hydrogen Energy 2007, 32, 2791–2802. (23) Funke, H.; Diaz, H.; Liang, X.; Carney, C. S.; Weimer, A. W.; Li, P. Int. J. Hydrogen Energy 2008, 33, 1127–1134. (24) Ernst, F. O.; Tricoli, A.; Pratsinis, S. E.; Steinfeld, A. AIChE J. 2006, 52, 3297–3303. (25) Melchior, T.; Piatkowski, N.; Steinfeld, A. Chem. Eng. Sci. 2009, 64, 1095–1101. (26) Abu Hamed, T.; Davidson, J. H.; Stolzenburg, M. J. Sol. Energy Eng. 2008, 130, 041010-1–041010-7. (27) Ehrensberger, K.; Palumbo, R.; Larson, C.; Steinfeld, A. Ind. Eng. Chem. Res. 1997, 36, 645–648E. (28) Yamasue, E.; Yamaguchi, H.; Nakaoku, H.; Okumura, H.; Ishihara, K. N. J. Mater. Sci. 2007, 42, 5196–5202. (29) Farghali, A. A.; Khedr, M. H.; Abdelkhalek, A. J. Mater. Proc. Technol. 2007, 181, 81–87. (30) Kato, H.; Kodama, M.; Tsuji, M.; Tamaura, Y. J. Mater. Sci. 1994, 29, 5689–5692. (31) Khedr, M. H.; Bahgat, M.; Nasr, M. I.; Sedeek, E. K. Colloids Surf., A 2007, 302, 517–524. (32) Khedr, M. H.; Farghali, A. A. Appl. Catal. B: Environ. 2005, 61, 219–226.
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metal oxides using pure streams of CO2 was experimentally demonstrated.27-41 In more recent studies, the thermodynamics of the solar-driven cycle were examined for CO2-splitting with Zn and FeO,42 and kinetic rate laws were determined for both reactions.43 A secondlaw (exergy) analysis for the net reaction CO2 = CO þ 0.5O2 indicated maximum solar-to-chemical energy conversion efficiencies of 39% and 29% for the Zn/ZnO and FeO/Fe3O4 cycles, respectively.42 This paper deals with the simultaneous reaction of H2O and CO2 with Zn or FeO to produce syngas. A thermodynamic analysis is performed to determine the equilibrium composition and energy requirements of the system over a selected range of temperatures, pressures, and stoichiometries and to guide the operating conditions for the experimental studies. An experimental thermogravimetric analysis is carried out to evaluate reaction kinetics and syngas quality. The thermodynamic and kinetic information establish the constraints to be imposed in the reactor design and process operation. 2. Thermodynamic Equilibrium Composition Thermodynamic equilibrium computations were carried out for the Zn/CO2/H2O and FeO/CO2/H2O systems, using the minimization of the free Gibbs energy function.44 Species with mole fractions of less than 10-5 were omitted from the figures. Figures 2 and 3 show the variation of the equilibrium compositions of the gas and solid phases, respectively, as a function of temperature at 1 bar for both Zn and FeO systems and at different stoichiometries: 2Zn þ H2O þ CO2 (Figures 2a and 3a), 3Zn þ H2O þ CO2 (Figures 2b and 3b), 6FeO þ H2O þ CO2 (Figures 2c and 3c), and 9FeO þ H2O þ CO2 (Figures 2d and 3d). The initial Zn or FeO to H2O/CO2 molar ratios were chosen to favor CO/H2 or C/CH4 production. For both Zn containing reactions, C and CH4 production are favored at lower temperatures and at higher initial Zn, although C is unlikely to be formed.43 For the 2Zn þ H2O þ CO2 system, the formation of H2 and CO is favored at intermediate and high temperatures, (33) Ma, L. Y.; Chen, L. S.; Chen, S. Y. J. Phys. Chem. Solids 2007a, 68, 1330–1335. (34) Ma, L. Y.; Chen, L. S.; Chen, S. Y. Mater. Chem. Phys. 2007b, 105, 122–126. (35) Tamaura, Y.; Tabata, M. Nature 1990, 346, 255–256. (36) Zhang, C. L.; Li, S.; Wang, L. J.; Wu, T. H.; Peng, S. Y. Mater. Chem. Phys. 2000, 62, 44–51. (37) Zhang, C. L.; Li, S.; Wu, T. H.; Peng, S. Y. Mater. Chem. Phys. 1999, 58, 139–145. (38) Zhang, C. L.; Liu, Z. Q.; Wu, T. H.; Yang, H. M.; Jiang, Y. Z.; Peng, S. Y. Mater. Chem. Phys. 1996, 44, 194–198. (39) Siegel, N. P.; Miller, J. E.; Diver, R. B.; Evans, L. R.; Branson, E. D.; Cooke, A. W. Presented at the 2007 annual AIChE meeting, November 3-9, Salt Lake City, UT, 2007 (40) Umeda, G. A.; Chueh, W. C.; Noailles, L.; Haile, S. M.; Dunn, B. S. Energy Environ. Sci. 2008, 1, 484–486. (41) Wu, S.; Zhang, X.; Gu, J. Wu Y.; Gao, J. Energy Fuels 2007, 21, 1827–1831. (42) Galvez, M. E.; Loutzenhiser, P.; Hischier, I.; Steinfeld, A. Energy Fuels 2008, 22, 3544–3550. (43) Loutzenhiser, P.; Galvez, M. E.; Hischier, I.; Stamatiou, A.; Frei, A.; Steinfeld, A. Energy Fuels 2009, 23, 2832–2839. (44) Roine, A. Outokumpu HSC Chemistry for Windows 5.0; Outokumpu Research Pori, Finland, 1997.
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Figure 2. Equilibrium composition of the gaseous products and solid carbon versus temperature at 1 bar for the systems (a) 2Zn þ H2O þ CO2, (b) 3Zn þ H2O þ CO2, (c) 6FeO þ H2O þ CO2, and (d) 9FeO þ H2O þ CO2.
respectively. H2 formation begins at above 500 K and peaks at about 1100 K, while CO formation begins at above 750 K and peaks at about 1250 K. Production of syngas is then thermodynamically favorable in the range 900-1050 K. The equilibrium behavior is analogous for the 3Zn þ H2O þ CO2 system, except that a significant amount of CH4 is observed at low temperatures. CO production is shifted to the right due to favored C production over a wider temperature range of 5001200 K and coincides with high quality syngas being shifted to higher temperatures. For both stoichiometries, complete conversion of Zn to ZnO is observed below 1000 K. Above 1000 K, the endothermic carbothermal reduction of ZnO becomes possible and Zn(g) is distilled in equilibrium. For the 6FeO þ H2O þ CO2 system, C formation is thermodynamically favorable over the temperature range 400-800 K. Small amounts of CH4 are formed up to
700 K. H2 appears at above 500 K and peaks at 900 K, while CO formation starts at above 800 K and becomes increasingly favorable above 1150 K. Between 400 and 500 K, the majority of the FeO is converted to Fe3O4 with residual amounts of Fe2O3. Above 500 K, nonstoichiometric wustite increases monotonically with temperature. From 800 to 1000 K, a small amount of FeO is reduced to Fe. For the 9FeO þ H2O þ CO2 system, the equilibrium composition is different at low temperatures. Between 400 and 450 K, there is an equimolar formation of CH4 and C. Above 450 K, CH4 decreases while C increases and peaks at 750 K. The H2, CO, and Fe containing products follow the same trend as the 6FeO þ H2O þ CO2 system. The enthalpy change of the pertinent reactions as a function of temperature is shown in Figure 4, with reactants at 298 K and products at equilibrium at the corresponding temperature. The ΔH of both systems increases with temperature mainly because of the increasing sensible
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Figure 3. Equilibrium compositions of the metal containing solid species versus temperature at 1 bar for the systems (a) 2Zn þ H2O þ CO2, (b) 3Zn þ H2O þ CO2, (c) 6FeO þ H2O þ CO2, and (d) 9FeO þ H2O þ CO2.
Figure 4. Enthalpy change of the reactions as a function of temperature for the chemical systems: (a) 2Zn þ H 2 O þ CO2 and 3Zn þ H 2 O þ CO2 and (b) 6FeO þ H2 O þ CO 2 and 9FeO þ H2 O þ CO 2 assuming reactants at 298 K and products at equilibrium at the given temperature.
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Figure 5. Reaction extent as a function of temperature at 0.1, 1, and 10 bar, for the systems 2Zn þ H2O þ CO2 (a, b) and 6FeO þ H2O þ CO2 (c, d).
heat absorbed. The 3Zn þ H2O þ CO2 reaction is more exothermal than the 2Zn þ H2O þ CO2 at lower temperatures due to the high exothermicity of C and CH4 formation. The change in the slope for the 3Zn þ H2O þ CO2 after 1000 K corresponds to lower conversions of CO2 and H2O to CH4 and C. The change of slope at same temperatures for 2Zn þ H2O þ CO2 system is also attributed to the decrease of C. Both systems turn endothermic at 1250 K. For the FeO/H2O/CO2 reactions, 9FeO þ H2O þ CO2 system is more exothermic than the 6FeO þ H2O þ CO2 system below 720 K, primarily because of higher CH4 formation. The high C formation also has a significant role, especially at the range of 600-800 K. Both reactions turn endothermic at 720 K, and above this point the system 6FeO þ H2O þ CO2 becomes the less endothermic. The reaction extents X are defined for the 2Zn þ H2O þ CO2 and 6FeO þ H2O þ CO2 systems with respect to C,
CO, and H2 in eq 3-5, respectively, as XC ¼
2neq C nin Zn
or
6neq C nin FeO
ð3Þ
XCO ¼
neq CO nin Zn
or
3neq CO nin FeO
ð4Þ
XH2 ¼
neq H2 nin Zn
or
3neq H2 nin FeO
ð5Þ
where nin and neq denote the initial and equilibrium amounts, respectively. XC, XCO, and XH2 as a function of temperature for pressures between 0.1 and 10 bar are given in Figure 5a,b for the 2Zn þ H2O þ CO2 system and in Figure 5c,d for the 6FeO þ H2O þ CO2 system.
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Table 1. Properties of Zn and FeO powders used in the TG experiments Material
dmean
SBET
Impurities
Titration
As: e0.1 mg/kg Cd: e500 mg/kg Cu: e50 mg/kg Fe: e50 mg/kg Pb: e500 mg/kg Sn: e10 mg/kg e5% free iron 99.9% Pure (based on TMI) 0.85, indicating a transition from an interfacecontrolled regime to a diffusion-controlled regime.43 For all experiments, XZn reached 0.85-0.95 after 45 min. The reaction for the CO2-H2O/FeO system commenced immediately after the reactive gas was introduced into the system. Initially, high reaction rates were observed over the first 5 min of the experiment, indicating
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Figure 7. Reaction extents (solid lines) and temperature (dashed lines) versus time for dynamic TG runs performed with (a) CO2-H2O/Zn and (b) CO2-H2O/FeO at various gaseous reactant concentrations.
Figure 8. Reaction extents (solid lines) and molar flow rates of CO (squares) and H2 (circles) versus time and temperature for the (a) CO2-H2O/Zn system and (b) CO2-H2O/FeO system. The molar flow fractions of the gas, R = 0.5.
rapid oxidation of the particle surface. Afterward, they remained relatively constant before transitioning to the diffusion-controlled regime after 45 min. For all experiments, XFeO > 0.8 after 50 min. For both systems, no direct correlations were identified between R and reaction rates. Figure 8a,b shows the product gas composition during exemplary dynamic runs for the CO2-H2O with Zn and FeO, respectively. XZn, XFeO, n_ H2, and n_ CO are given as functions of t and T at R = 0.5. For the CO2-H2O/Zn system, the production of CO and H2 started at about 780 K. The molar flow rate of H2 peaked at 1000 K whereas the one for CO peaked at 1060 K, which is consistent with the thermodynamic calculations favoring CO at higher temperatures. The amount of H2 produced was much higher than the corresponding CO, indicating a kinetically favored reaction of Zn with H2O rather than with CO2. The profile of gas production is remarkably different for
the CO2-H2O/FeO system, which took place at higher T. A high peak of H2 was observed at the beginning of the reaction which corresponded to the initial fast mass gain measured by the TG. This peak indicated that H2O was preferentially absorbed on the surface of FeO particles. After 1000 K, H2 production stabilized and remained constant until 1300 K where it decreased. The CO production reached its highest value at 950 K and remained almost constant until 1300 K when it started to decrease. It should be noted that other intermediate reactions might also take place at the systems examined. More specifically, the water-gas shift reaction (WGSR) may occur, altering the composition of the product gases as they travel downstream to the GC. The thermodynamics predict that below 1100 K the WGSR takes place, shifting the equilibrium toward H2 production, whereas above 1100 K the reverse WGSR is thermodynamically favored, shifting the equilibrium toward CO production.
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Figure 9. Reaction extents (solid lines) and the accumulated normalized amounts of CO (squares), H2 (circles), and H2 þ CO (diamonds) versus time for the (a) CO2-H2O/Zn system and (b) CO2-H2O/FeO system. The molar flow fractions of the gas, R = 0.2.
Although thermodynamically favorable, no C or CH4 formation was observed throughout the experiments, as verified by GC and mass balances. XZn and XFeO and the normalized amounts of accumulated H2, CO, and H2 þ CO are given for CO2-H2O/Zn and CO2-H2O/FeO systems in Figure 9a,b, respectively, as a function of t with R = 0.2. The normalized amounts of accumulated H2 þ CO from the GC predict slightly lower reaction extents compared with the XZn and higher reaction extents compared XFeO. For the CO2-H2O/FeO system, the CO and H2 concentrations were below the detection limit in the diffusion-controlled regimes. The solid products were only ZnO and Fe3O4, as verified by X-ray diffraction analysis. This can be attributed to kinetic hindrances as well as to the fact that, through the continuous feeding of reactive gas and removal of the gaseous products, the experimentally investigated system was not allowed to reach thermodynamic equilibrium. Both systems exhibited a strong dependency on R, which can be seen in Figure 10. Integrated values of γ over time were used to compare the experiments. A linear dependency between γ and R for both reactions was observed. These correlations reveal the possibility of adjusting the syngas composition by varying the reactant flows. 4. Summary and Conclusions Two-step solar thermochemical cycles were examined to reduce simultaneously CO2 and H2O to syngas using Zn/ZnO and FeO/Fe3O4 redox pairs. Thermodynamic equilibrium computations revealed that the formation of C is favored at below 1000 K and high quality syngas can be produced for both systems at above 900 K. The two reactions were also investigated experimentally by thermogravimetry and gas chromatography. During the interface controlled regime for both Zn and FeO, H2O exhibited higher reaction rates with the solids compared to CO2, indicative of preferential absorption of H2O on
Figure 10. Product molar ratio γ as a function of reactive gas molar ratio R for the CO2-H2O/Zn (circles) and CO2-H2O/FeO (squares) systems.
the particle surfaces. A strong dependency between the H2O/CO2 molar ratio of the input gases and the H2/CO molar ratio of the product gases was shown over the temperature ranges investigated. These results provide guidance for the chemical reactor design and operation. The proposed cycles offer a viable alternative to the sequestration of captured CO2. Acknowledgment. This work has been partially supported by the Baugarten Foundation and the Swiss Federal Office of Energy. The authors thank G. Illari and A. Frei at PSI for their support with the experimental campaign and M.E. G alvez from the Instituto de Carboquimic in Spain for her help with the TG runs.
Nomenclature m = mass M = molecular weight Δm = mass change n = moles
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n_ = molar flow rate p = pressure t = time T = temperature X = reaction extent ΔH = enthalpy change
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R = gaseous reactant molar flow fraction γ = gaseous product molar flow fraction Superscript eq = equilibrium in = initial
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