Hydrogen Production by Three-Step Solar Thermochemical Cycles

Energy & Fuels 2007, 21, 2919-2928

2919

Hydrogen Production by Three-Step Solar Thermochemical Cycles Using Hydroxides and Metal Oxide Systems Patrice Charvin,*,† Ste´phane Abanades,† Florent Lemort,‡ and Gilles Flamant† CNRSsProcesses, Materials and Solar Energy Laboratory, 7 Rue du Four Solaire, 66120 Font Romeu, France, and Commissariat a` l’Energie Atomique, BP 17171, 30 207 Bagnols-sur-Ce` ze Cedex, France ReceiVed March 26, 2007. ReVised Manuscript ReceiVed June 14, 2007

This paper presents a thermodynamic and experimental study of three-step thermochemical cycles for hydrogen production involving hydroxides (NaOH and KOH). Solar concentrated energy was successfully used to reduce manganese, cobalt, and iron oxides into lower valence metal oxides, MnO, CoO, Fe3O4, and FeO, in the temperature range of 1300-1600 °C. In the reaction with NaOH and KOH, MnO and CoO were stable and did not produce hydrogen at 750 °C even in a strong oxidizing media, whereas the iron oxides FeO and Fe3O4 were able to generate hydrogen. For the NaOH activation reaction, the final chemical conversion rate was 28% at about 400 °C with FeO particles in the range of 30-50 µm, and a passivating layer was observed, which reduced the H2 production rate when the particle size increased. The reaction between Fe3O4 particles and NaOH reached a final conversion higher than 70% after 7 min of the reaction for particle sizes in the range of 30-125 µm. In addition, the reaction between Fe3O4 and KOH producing hydrogen was nearly complete. Although the three-step cycle based on FeO appears attractive in terms of theoretical productivity (156 mLH2 g-1 of FeO assuming a complete reaction) and energy efficiency (41.3%), it requires a hightemperature reduction reaction and a small particle size for the H2-production reaction. Finally, the comparison of iron oxide cycles highlights the high potential of the three-step cycle based on the Fe2O3/Fe3O4 pair, taking into account experimental chemical conversions (37 mLH2 g-1 of Fe3O4 for a 75% chemical conversion).

1. Introduction Solutions are needed to reduce greenhouse gas emissions because of fossil fuel combustion in the transportation sector. Hydrogen is a promising energy carrier for the substitution of fossil fuels in the next decades. This sustainable fuel has no impact on the environment, if produced from, e.g., water and solar energy.1 Solar energy is renewable, and water consumed for hydrogen production is given back to the environment when H2 is processed in a fuel cell. Large-scale production processes must be environmentally friendly and need reasonable investments for a low cost of hydrogen. Short thermochemical cycles using high-temperature solar energy potentially offer higher theoretical energy conversions than electrolysis, which requires heat-electricity conversion before electrolyzing water (23% global efficiency for a 33% nuclear heat-electricity efficiency). In the case of photovoltaic electricity, the global electrolysis efficiency decreases to about 14%, assuming an electricity production efficiency of 20%. Thermochemical water-splitting consists of the conversion of water into hydrogen and oxygen by a series of endo- and exothermic chemical reactions. All of the intermediate chemicals are recycled within the process. A previous selection of promising thermochemical cycles2 showed that two- and threestep thermochemical cycles using metal oxides are very attrac* To whom correspondence should be addressed. Fax: +33-4-68-3029-40. E-mail: [email protected]. † CNRSsProcesses, Materials and Solar Energy Laboratory. ‡ Commissariat a ` l’Energie Atomique. (1) Koroneos, C.; Dompros, A.; Rombas, G.; Moussiopoulos, N. Int. J. Hydrogen Energy 2004, 29, 1443-1450. (2) Abanades, S.; Charvin, P.; Flamant, G.; Neveu, P. Energy 2006, 31, 2805-2822.

tive for coupling with solar energy. The high reduction temperature (above 900 °C) does not correspond to nuclear requirements3 but is suitable for solar concentrating technologies. Two-step cycles proceed with the reduction of an oxide or sulfate in an endothermic reaction at high temperature. The reduced oxide reacts with water at a low temperature to regenerate the first oxide and produce hydrogen

MOox f MOred + 1/2O2 (T > 1300 °C)

(1)

MOred + H2O f MOox + H2 (T < 1000 °C)

(2)

The ZnO/Zn redox pair proposed by Bilgen4 has been studied thoroughly,5-7 as have other metal oxide redox pairs, such as Fe3O4/FeO.8 An experimental study9 of both reactions gave details about the iron oxide cycle and indicated the optimal conditions and conversion rates. Design aspects concerning a solar reactor are discussed in ref 10. An alternative to a pure iron oxide cycle is a process based on using mixed oxides with a FeO matrix and the addition of small amounts of nickel, (3) Brown, L. C.; Besenbruch, G. E.; Lentsch, R. D.; Schultz, K. R.; Funk, J. F.; Pickard, P. S.; et al. GA-A24285. Nuclear Energy Research Initiative Program for the U.S. Department of Energy, 2003. (4) Bilgen, E.; Bilgen, C. Int. J. Hydrogen Energy 1981, 7, 637-644. (5) Mo¨ller, S.; Palumbo, R. Chem. Eng. Sci. 2001, 56, 4505-4515. (6) Steinfeld, A. Int. J. Hydrogen Energy 2002, 27, 611-619. (7) Wegner, K.; Ly, H. C.; Weiss, R. J.; Pratsinis, S. E.; Steinfeld, A. Int. J. Hydrogen Energy 2006, 31, 55-61. (8) Nakamura, T. Sol. Energy 1976, 19, 467-475. (9) Charvin, P.; Abanades, S.; Flamant, G.; Neveu, P.; Lemort, F. Energy 2007, 32, 1124-1133. (10) Steinfeld, A.; Sanders, S.; Palumbo, R. Sol. Energy 1999, 65, 4353.

10.1021/ef0701485 CCC: $37.00 © 2007 American Chemical Society Published on Web 07/24/2007

2920 Energy & Fuels, Vol. 21, No. 5, 2007

CharVin et al.

Figure 1. Equilibrium composition results of the Mn/O system.

manganese,11 cobalt, or zinc12 to lower the reduction temperature and produce a reduced oxide that is still reactive with water. Co3O4/CoO and MnO2/MnO redox pairs were also proposed with the objective of lowering the maximum cycle temperature. Unfortunately, water-splitting reactions with CoO and MnO are not thermodynamically favorable, with conversion rates lower than 1%.13 The substitution of water with a better oxidizing compound, such as sodium hydroxide, may enhance the oxidation of the reduced metal oxide during the hydrogen production reaction. In this case, the product rapidly splits water in the last step at nearly ambient temperature. Such three-step cycles proposed by Perkins and Weimer14 and Sturzenegger et al.15 in the case of manganese oxide may be possible for all reduced oxides

MOox f MOred + 1/2O2 (T > 1300 °C) MOred + 2NaOH f Na2O‚MOox + H2 (T < 700 °C)

(3) (4)

Na2O‚MOox + H2O f MOox + 2NaOH (T ) 100 °C) (5) For example, the use of Fe3O4 as a reduced oxide in a threestep cycle was considered in ref 16. Little information about the reaction temperature and optimal conditions is available on Fe3O4 and Mn2O3 reduction17 and reoxidation.18 Moreover, the Co3O4/CoO cycle has never been studied experimentally before. The theoretical hydrogen productivity of Fe2O3/Fe3O4, Co3O4/CoO, Fe2O3/FeO, and Mn2O3/ (11) Tamaura, Y.; Kojima, M.; Sano, T.; Ueda, Y.; Hasegawa, N.; Tsuji, M. Int. J. Hydrogen Energy 1998, 23, 1185-1191. (12) Aoki, H.; Kaneko, H.; Hasegawa, N.; Ishihara, H.; Suzuki, A.; Tamaura, Y. Solid State Ionics 2004, 192, 113-116. (13) Lundberg, M. Int. J. Hydrogen Energy 1993, 18, 369-376. (14) Perkins, C.; Weimer, A. W. Int. J. Hydrogen Energy 2004, 29, 1587-1599. (15) Sturzenegger, M.; Nuësch, P. Energy 1999, 24, 959-970. (16) Bamberger, C. E.; Richardson, D. M.; Grimes, W. R. U.S. Patent 3,929,979, December 30, 1975. (17) Sibieude, F.; Ducarroir, M.; Tofighi, A.; Ambriz, J. Int. J. Hydrogen Energy 1982, 7, 79-88. (18) Sturzenegger, M.; Ganz, J.; Nuësch, P.; Schelling, T. J. Phys. IV 1999, 9, 331-335.

MnO cycles is 4.3, 8.9, 13.9, and 14.1 mgH2 g-1 of reduced oxide, respectively. In this study, the solar reduction of Fe2O3, Mn2O3, and Co3O4 is presented. In addition, the reaction of hydroxide activation producing hydrogen from CoO, MnO, FeO, and Fe3O4 was performed in test tubes to determine the reaction temperature and chemical conversion rate. The quantitative analysis of H2 produced was realized continuously during the activation step with NaOH and KOH for different metal oxides and particle sizes. The hydrolysis reaction regenerating the initial metal oxide and hydroxide was also investigated. 2. Thermodynamic Analysis The minimization of Gibbs free enthalpy was realized by the HSC Chemistry software19 (equilibrium composition module) to predict stable species quantities versus temperature. For given operating conditions (T, P, and composition), the theoretical reaction conversion and possible secondary reactions were determined. Results obtained for the different redox pairs (Fe2O3/ Fe3O4, Fe2O3/FeO, Mn2O3/MnO, and Co3O4/CoO) were compared with experimental ones. 2.1. High-Temperature Reduction Reactions. The quantities of stable products at equilibrium were determined for different temperatures. Figure 1 illustrates the case of the Mn2O3/MnO pair. The conversion rate of the reduction reaction was estimated. It is defined as the ratio of the amount of product obtained to the amount of product that could be obtained if the reaction was complete. The reaction temperatures and chemical conversions obtained from thermodynamics for the redox pairs mentioned above are presented in Table 1. The lower the pressure, the lower the reaction temperature. A decrease in the reaction temperature represents an important energy saving, although an additional work input is necessary to reduce the pressure by vacuum pumps. In the case of manganese oxide, the energy required to heat Mn2O3 and produce MnO is reduced from 527.89 kJ/mol at 2000 °C to 366.89 kJ/mol at 1470 °C. Accounting for the heating of the (19) Roine, A. HSC Chemistry 5.11; Outokumpu Research Oy, Pori, Finland, 2002.

Three-Step Solar Thermochemical Cycles

Energy & Fuels, Vol. 21, No. 5, 2007 2921

Table 1. Reaction Degree of Conversion Given by Thermodynamics (N2 Atmosphere) reactant

product

pressure (bar)

reaction temperature (°C)

conversion (%)

Fe2O3 Fe2O3 Fe2O3 Fe2O3 Fe2O3 Mn2O3 Mn2O3 Mn2O3 Mn2O3 Co3O4

Fe3O4 FeO FeO FeO FeO MnO MnO MnO MnO CoO

1 1 0.1 0.01 0.001 1 0.1 0.01 0.001 1

1250 2100 1900 1700 1600 2000 1800 1620 1470 1300

60 98 93 92 90 96 95 95 95 100

3. Experimental Study 3.1. High-Temperature Reductions. 3.1.1. Reaction Temperature Determination. Reduction temperatures were determined precisely by thermogravimetric analysis (TGA; Setsys device of SETARAM). Oxygen released during the reaction induced an important weight loss, which was used to identify the reaction temperature (Figure 5). At atmospheric pressure under an inert atmosphere, the reaction temperatures obtained were

Fe2O3/Fe3O4 Treaction ) 1300 °C

reactants for the low-temperature exothermal reactions (177.38 kJ/mol), the theoretical energy efficiency based on the HHV of hydrogen is equal to 52.5% at 1470 °C and 38% at 2000 °C. The effect of the temperature on solar reactor exergy efficiency is even more important because of a reduction of the radiative losses (T4 dependency). Thermodynamics shows that the synthesis of iron(II), manganese(II), and cobalt(II) oxides from their higher valence states by solar reduction at atmospheric pressure is possible at temperatures below 2100, 2000, and 1300 °C, respectively. 2.2. Activation Reaction with Sodium Hydroxide. The reaction between a reduced oxide and water is thermodynamically possible if the Gibbs free energy of the metal oxide redox pair is greater than the ∆G of water reduction. Figure 2 shows that water-splitting is only possible with FeO below 800 °C. MnO, CoO, and Fe3O4 are not able to split water spontaneously, which means that a work input is necessary to run the reaction (∆G > 0). The alternative consists of using a better oxidizing compound, such as sodium or potassium hydroxide, to conserve a pure thermal cycle. As shown in Figure 2, the oxidation of MnO and Fe3O4 is possible with sodium hydroxide for temperatures below 700 °C. Inversely, CoO remains unreactive with NaOH. The reaction with NaOH (or KOH) produces hydrogen and a mixed oxide. According to thermodynamics, a complete conversion can be obtained at low temperatures for FeO (Figure 3) with the formation of a mixed oxide NaFeO2, which can also be written as Na2O‚Fe2O3. In the case of Fe3O4, hydrogen can also be produced at a low temperature, but the H2-generation reaction competes with the disproportionation of Fe3O4 (Fe3O4 f FeO + Fe2O3). An excess of NaOH in the initial mixture (10 mol of NaOH/mol of Fe3O4) is enough to favor the H2generation reaction, which becomes complete in the range of 100-1000 °C (Figure 4). The same study could not be conducted for Mn(II) and Co(II) oxides because thermodynamic data of the mixed oxides are not available. The results predicted by thermodynamics can be modified because of kinetic limitations, and the minimum reaction temperature must be determined from experiments. 2.3. Sodium Hydroxide Regeneration. Thermochemical cycles impose the requirement for regeneration of all reactants and the consumption of water. A third reaction must be carried out between the mixed oxide and water. In the case of iron oxide, it can be written as

Na2O‚Fe2O3 + H2O f Fe2O3 + 2NaOH

which requires an additional step of separation, e.g., by evaporating water.

(6)

The Gibbs free energy of the reverse reaction is positive (Figure 2), which means that sodium hydroxide regeneration is thermodynamically feasible at ambient temperature. If an excess of water is used, the NaOH is obtained as an aqueous solution,

Fe3O4/FeO Treaction ) TmFe3O4 ) 1600 °C Mn3O4/MnO Treaction ) 1450 °C Co3O4/CoO Treaction ) 900 °C 3.1.2. Solar Experimental Setup. Reduction of metal oxides requires temperatures above 900 °C, which can be obtained from concentrated solar energy. A solar furnace composed of a reflector (flat heliostat) and a parabolic mirror (1.5 m in diameter, peak flux density of 16 MW/m2) provided solar heat at a high temperature to the sample. Rough control of the heating can be achieved by moving the solid sample vertically with respect to the focal plane. A quick withdrawal of the sample from the focus provides a rapid quenching (about 100 °C/s).20 Commercial powders of pure Fe2O3 (99.9%, size of 100 µm), monohydrated MnO2‚H2O (99.9%), and nickel-free Co3O4 (Co3O4, 99.45%; sulfates, 0.5%) were directly heated or preliminary compacted into pellets. The concentrated solar radiation was directed to a water-cooled holder, where a sample of metal oxide (about 1 g) had been placed. In some experiments, a glass vessel was added around the sample to control the atmosphere (Figure 6). An inert gas inlet (nitrogen) and a vacuum pump connected to the gas outlet were used to work under an inert atmosphere and/or at reduced pressure. The continuous flow of inert gas was used to sweep oxygen evolved during the reduction reaction. Thus, reoxidation of the reduced metal oxide with oxygen was not favored. After the experiment, the solid sample was crushed into a fine powder and analyzed with X-ray diffraction (XRD; Philips X’Pert analyzer) using a Cu KR radiation source. 3.1.3. Reduction of Iron(III) Oxide. Reduction of iron(III) oxide into FeO was described previously.9,17,21 The kinetics of the reduction reaction were slow, which can be explained by the difference between the reaction temperature and the Fe3O4 melting point. At the melting point, the Gibbs free energy between the solid and liquid (eq 7) is equal to 0. Then, the evolution of the temperature versus pressure is given by eq 8.

∆Gsfl ) ∆Hsfl - T∆Ssfl + P∆Vsfl ) 0 Tsfl ) -

∆Hsfl ∆Vsfl +P sfl ∆S ∆Ssfl

(7) (8)

Pressure has a very low influence on the melting point because of the low volume variation between a solid and a liquid phase. (20) Tofighi, A. Ph.D. Thesis. Institut National Polytechnique de Toulouse, France, 1982 (in French). (21) Tofighi, A.; Sibieude, F. Int. J. Hydrogen Energy 1984, 9, 293296.


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Figure 2. Gibbs free energy of reduction reactions.

Figure 3. Equilibrium composition results of the Fe/O/Na/H system with initial conditions of 2 mol of FeO, 2 mol of NaOH, 100 mol of N2, and a total pressure of 1 bar.

Thus, the melting point can be considered constant as a function of the pressure. Inversely, pressure has a great influence on the reaction temperature (Table 1). For a pressure of 1 bar, the theoretical reaction temperature given by thermodynamics in Table 1 (2100 °C) is higher than the Fe3O4 melting point (1597 °C), which was the working temperature. Indeed, the working temperature is just above the melting temperature because of good heat exchange between the liquid sample and the water-cooled support. The higher the difference between the working and theoretical temperatures, the lower the reaction rate. Thus, the reaction rate increased with a pressure reduction. Experimentally, the time elapsed for 95% reaction completion was more than 5 min at atmospheric pressure and about 2 min at 0.1 bar. Temperature measurements, realized with a “solar blind” pyrometer at 5.2 µm,22 corrected by calculated emissivities (with

the melting point as the reference temperature, 5.2 µmFe3O4 (s) ) 0.79, 5.2 µm Fe3O4 (l) ) 0.69, and 5.2 µm FeO (l) ) 0.56) showed two successive stabilization periods (Figure 7). The first one is just above the Fe3O4 melting point (TmFe3O4 ) 1597 °C), and the second one corresponds to the FeO melting point (TmFeO ) 1377 °C). It confirmed that the formation of FeO started after Fe3O4 melting. The decrease of the temperature below the FeO melting point may be due to the endothermic reaction and the change of material properties during the reaction (both radiative and thermodynamic properties change from Fe3O4 to FeO). For the transition from Fe2O3 into Fe3O4, the reaction temperature was lower than for FeO synthesis and reduction (22) Hernandez, D.; Olalde, G.; Gineste, J. M.; Gueymard, C. J. Sol. Energy Eng. 2004, 126, 645-653.



Figure 4. Equilibrium composition results of the Fe/O/Na/H system with initial conditions of 1 mol of Fe3O4, 10 mol of NaOH, 100 mol of N2, and a total pressure of 1 bar.

Figure 5. TGA and differential thermal analysis (DTA) of Fe2O3 under an Ar atmosphere.

kinetics were very rapid as previously indicated by Tofighi.20 This first reduction at the solid state can be realized under air at about 1300 °C. 3.1.4. Reduction of MnO2. Manganese(IV) oxide was used as a raw material to test the thermal production of pure MnO. After MnO reoxydation with sodium hydroxide and water, the reduction (second cycle) involves Mn(III) oxide (Mn2O3) rather than MnO2. The complete reduction of MnO2 into MnO was obtained after 1 min of heating under a N2 atmosphere after melting the sample (TmMn3O4 ) 1562 °C). The rapid kinetics of reduction meant that the sample was heated above the reaction temperature (1450 °C). The reaction cannot be conducted under air because the quenching rate is not high enough to prevent reoxidation with oxygen,16 and as a result, a mixture of Mn3O4 and MnO is produced after 5 min of the reaction.

3.1.5. Reduction of Co3O4. The reduction of Co3O4 into CoO was investigated. The sample melting point was reached after 10 s. CoO melting occurs experimentally at 1830 °C, which suggests a large overheating of the sample, because thermodynamics predicts a reaction temperature around 1300 °C. The reaction kinetics are rapid, and complete conversion is obtained after less than 10 s. In contrast to iron and manganese oxide reduction, an inert atmosphere is not required for the reduction of Co3O4. The reoxidation kinetics are too slow to convert CoO during quenching (100 °C s-1). Thus, this rapid reduction under air is easy to implement in a large-scale process. The solar reactor does not require a quartz window, but the low reactivity of CoO with oxygen during quenching suggests that hydrolysis will be difficult. 3.2. Activation Reaction with Sodium Hydroxide Producing Hydrogen. The solar-reduced metal oxides were prepared


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Figure 6. Experimental setup at the focus of a solar furnace.

Figure 7. Temperature measurements during the heating of a Fe2O3 pellet.

in large quantities to test the sodium hydroxide activation reaction producing hydrogen (eq 4). The solidified pure sample of solar-reduced oxide was milled into a fine powder to increase its specific surface area. 3.2.1. Experimental Setup. The activation reaction involved solar-reduced solid oxide and sodium hydroxide (Aldrich, 97%, pellets). The two reactants were introduced together in an Inconel test tube at the beginning of the experiment. Sodium hydroxide was used as the reactant and solvent (liquid phase), and it was introduced in large excess (6 mol for 1 mol of metal oxide). The test tube was sealed and connected to an inert gas (Ar) feeder, which swept the produced gas to the reactor outlet (Figure 8). After a purge of the setup with argon to eliminate air, the test tube was placed in an electrical tubular furnace preliminarily heated at the working temperature. The temperature increased progressively in the test tube to reach the melting point of sodium hydroxide (323 °C), at which time a temperature plateau occurred. Once NaOH completely melted, the temperature increased again up to the working temperature. A liquid phase

of sodium hydroxide was obtained with reduced oxide particles in suspension. The flow of argon (0.18 NL/min) carried hydrogen produced during the reaction to the analytical system. It included a gas treatment device composed of a bubbler (water trap) and a gas dryer connected to the gas outlet to provide a dry mixture of argon and hydrogen to the continuous gas analyzer (catharometer ARELCO Catarc 10P; detection limit, 100 ppm; precision, 1% of full scale). The molar flow rate of hydrogen (FH2) was evaluated from the mole fraction of H2 (yH2) measured by the catharometer

F H2 )

yH 2 1 - yH2

FAr

(9)

The total volume of hydrogen produced during an experiment was calculated by integrating the production curve from the starting time (t0) to the ending time of the experimental run (tf)


V H2 )

∫t t FH

RT P

f

0

2

dt


(10)

where FH2 is the molar flow rate of H2 and VH2 is the volume of H2 produced at normal conditions (T ) 0 °C, and P ) 1.013 bar). The corrosiveness of reactants is an important issue that may be encountered during experiments. Thus, a metal alloy (Inconel) was used for the test-tube material at operating temperatures below 700 °C. 3.2.2. Results and Discussion: Hydrogen Production from MnO and CoO. For both reduced Mn and Co oxides, no hydrogen production was observed after the melting of sodium hydroxide. The temperature increased up to 800 °C in 5 min, but a problem of material corrosion was observed after a few minutes. Above 700 °C, the metal alloy was oxidized by the corrosive compounds involved in the reaction (Na2O and NaOH),23 producing hydrogen, which disturbed measurements. Because of these material resistance problems, it was impossible to maintain the reactants at 700 °C for 75 min, as suggested by Sturzenegger et al.18 for MnO reoxidation at atmospheric pressure. The resistance of reactor materials to corrosive chemicals at high temperatures is an important issue that must be taken into account in the choice of a thermochemical cycle. Such cycles based on Mn or Co oxides require expensive materials, such as platinum, silver, or nickel, which makes the cycles not viable. Hydrogen Production from FeO. The production of hydrogen was measured at the beginning of the sodium hydroxide melting for FeO produced at the focus of a solar furnace, as illustrated in Figure 9. At this time, the main change occurring in the reactor was the contact between reactants enhanced by the appearance of liquid sodium hydroxide. An increase in hydrogen production was observed at the end of sodium hydroxide melting (at time t ) 280 s) because the temperature increased again. The reaction kinetics were rapid when the temperature was higher than 300 °C, and hydrogen production became negligible after 15 min. Conversion rates were low and increased significantly with a particle size decrease (Table 2). For a given volume, a relation of inverse proportionality can be demonstrated between the total surface area of a sample and the particle diameter (eq 11).

3 S 4πR2 3 ) ) wS) V V 4 3 R R πR 3

(11)

The same relation was observed for chemical conversion versus the particle diameter (Figure 10), which proved that sodium hydroxide reacted only at the surface of the FeO particle. As already demonstrated for a two-step cycle,24,25 the formation of a nonpermeable oxidized layer stops the liquid-solid reaction. High conversion rates can be obtained with a fine ball milling to reduce the particle diameter and increase their surface area. An XRD analysis (Figure 11) confirmed the formation of the targeted product NaFeO2 at the bottom of the test tube after elimination of sodium hydroxide by dissolution in water. Peaks of FeO were still visible in the spectrum because the reaction was not complete. (23) Kolchakov, J.; Tzvetkoff, T.; Bojinov, M. Appl. Surf. Sci. 2005, 249, 162-175. (24) Weidenkaff, A.; Nuësch, P.; Wokaun, A.; Reller, A. Solid States Ionics 1997, 101-103, 915-922. (25) Ehrensberger, K.; Frei, A.; Kuhn, P.; Oswald, H. R.; Hug, P. Solid State Ionics 1995, 78, 151-160.

Figure 8. Experimental setup used for NaOH and KOH activation reactions.

The same activation reaction was conducted with potassium hydroxide. Hydrogen was produced at lower temperatures (250 °C) with rapid kinetics, but final conversion rates were similar. Hydrogen Production from Fe3O4. Hydrogen was generated as soon as the melting of sodium hydroxide occurred at 300 °C, and the production profile was similar to that obtained for the reaction with FeO. The reaction rate and maximum of hydrogen detected in the outlet gas were slightly lower than for FeO samples milled at the same size. Fe3O4 is composed of one atom of FeII and two atoms of FeIII. Thus, hydrogen production from Fe3O4 is theoretically limited to one-third that from FeO. Integration of production curves gave similar or higher amounts of hydrogen evolved for the same mass of solid reactant (Table 3). Then, Fe3O4 chemical conversions were 3 times higher than FeO ones. No significant difference could be detected in the final chemical conversion for different Fe3O4 particle sizes (Table 3). However, a particle size decrease led to an increase of the solid/liquid interface area, which results in higher reaction rates, as shown in Figure 12. Unlike with FeO, Fe3O4 reoxidation was not interrupted by a slow diffusion of NaOH through a passivating layer of oxide at the surface of particles. A final chemical conversion of about 75% was reached. Reaction kinetics could not be precisely determined because reactants were not introduced at a constant temperature. Sodium hydroxide was replaced by potassium hydroxide to compare reactivity of both oxidizing agents. Experimental results, in agreement with ref 16, showed a starting reaction temperature of 270 °C, similar to sodium hydroxide, but with an enhanced rate of reaction. The chemical conversion for different particle sizes reached 98% with potassium hydroxide. 3.3. Hydrolysis of the Mixed Oxide and Sodium Hydroxide Regeneration. A thermochemical cycle is based on the regeneration of all reactants except water within the cycle. Although the key reaction in the three-step cycle is the sodium hydroxide activation, hydrolysis of the product (NaFeO2) must be demonstrated to close the cycle and regenerate metal oxide and sodium hydroxide. Iron Oxide Cycles. The activation reaction with NaOH produced a mixed oxide (NaFeO2), which must be hydrolyzed


CharVin et al.

Figure 9. Mole fraction of hydrogen and temperature versus the time recorded during a reaction between 2 g of FeO (30 < dp < 50 µm) and NaOH (FAr ) 0.18 NL/min).

Figure 10. Evolution of the reaction chemical conversion versus the diameter of the particles. Table 2. Hydrogen Production during the Reaction between NaOH and FeO experiment

mass (g)

size (µm)

H2 volume (mL)

chemical conversion (%)

reaction time (min)

mmol of H2/ g of FeO

1 2 3 4 5 6 7

3 1.42 1.5 3 2 1.31 1.5

150 < x < 315 80 < x < 100 50 < x < 80 30 < x < 50 30 < x < 50 5 < x < 30 10 < x < 20

21.6 30.8 45.7 114.3 85.9 104.8 125

4.7 13.9 19.5 24.4 27.5 51.2 53.5

13 14 12.3 15 17 16 13

0.32 0.97 1.36 1.7 1.92 3.57 3.72

to regenerate Fe2O3 and sodium hydroxide. At the end of the previous step, the cooling of products led to the solidification of the solvent (NaOH). This solid was dissolved with water to separate sodium hydroxide in excess from the mixed oxide. After decantation and drying of the solid product, NaFeO2 was detected, which showed that the hydrolysis reaction (eq 6) regenerating NaOH and Fe2O3 did not occur at room temperature. The mixture of water and NaFeO2 was heated to the water boiling point to carry out the hydrolysis reaction (eq 6). XRD analysis of the hydrolyzed sample revealed that pure Fe2O3 was obtained, as already observed in ref 26, allowing its reuse in a second cycle. Consequently, the three-step iron oxide thermochemical cycle was completely demonstrated. Concerning the

recovery of NaOH from its aqueous solution, a step of drying for removing water was necessary. Then, the additional required energy input should be included when dealing with energyefficiency considerations. 4. Discussion 4.1. Chemistry. The reactivity with sodium hydroxide strongly depends upon the metal oxide used. For cobalt oxide, the hydrogen production is impossible at atmospheric pressure. CoO is very stable at medium temperature even when mixed (26) Blesa, M. C.; Moran, E.; Menendez, M.; Tornero, J. D.; Torron, C. Mater. Res. Bull. 1993, 28, 837-847.



Figure 11. XRD patterns of the solid sample after NaOH activation.

Figure 12. Evolution of conversion rates for NaOH and KOH activation reactions with Fe3O4. Table 3. Hydrogen Production during the Reaction between NaOH (or KOH) and Fe3O4 experiment

hydroxide

mass (g)

size (µm)

H2 volume (mL)

chemical conversion (%)

reaction time (min)

mmol of H2/ g of Fe3O4

1 2 3 4

NaOH NaOH NaOH KOH

1.5 1 1 1.5

50 < x < 80 100 < x < 125 30 < x < 50 50 < x < 80

55.4 37.4 35.2 71.3

76 77 73 98

15 15 7 7

1.65 1.67 1.57 2.12

with a strong oxidizing compound (NaOH). Thus, the threestep cycle with cobalt oxide is not feasible below 750 °C. Concerning manganese oxide, previous experimental results18 demonstrated the NaOH activation reaction at 650 °C. However, a dynamic vacuum was required to obtain reasonable reaction rates at this temperature. The main problem of the manganese oxide three-step cycle is the resistance of reactor walls under very corrosive conditions. A series of metallic alloys were tested, and each one was oxidized with sodium hydroxide generating hydrogen. The previous study18 was carried out in Ag and Cu crucibles coupled with a TGA measurement. These expensive materials increase investments, thereby increasing the cost of hydrogen produced. The results of the FeO reaction with sodium hydroxide are similar to the reaction with steam (direct hydrolysis in a two-

step cycle). The two- and three-step cycles have the reduction reaction, which produces FeO in the same conditions of temperature and pressure in common. The comparison of both cycles must focus on the hydrogen-generation step. The formation of an oxide layer, which stops the reaction, is common and requires a fine ball milling to reduce the particle size and increase the particle surface area. The activation with sodium hydroxide is more rapid than direct hydrolysis, and it occurs at a lower temperature (400 °C instead of 600 °C for direct hydrolysis).9 From the chemical point of view, the amount of FeO necessary to produce 1 mol of hydrogen is 2 mol in the three-step cycle and 3 mol in the two-step cycle. For the same particle size, the chemical conversion of the NaOH activation reaction (28%) is lower than the chemical conversion of direct hydrolysis (54% in ref 9). This low conversion at low temper-


CharVin et al.

Table 4. Comparison Results of Iron Oxide Cycles criterion reduction temperature (°C) H2-generation temperature (°C) productivity (mg of H2/ g of iron oxide) energy efficiency (HHV H2) (%)

Fe3O4/FeO (two step)

Fe2O3/FeO (three step)

Fe2O3/Fe3O4 (three step)

1600

1600

1300

600

400

400

9.3

13.9

4.3

33.4

41.3

25.3

atures can be linked to iron ion diffusion (temperaturedependent) and nonstoichiometry in FeO.27 Then, the amount of hydrogen produced per gram of FeO is quite similar in both cycles (50 mL/g of FeO for the particle size in the range of 30-50 µm). Although Fe3O4 and MnO have the same Gibbs free energy, Fe3O4 reacts rapidly with sodium hydroxide at low temperature. Unlike FeO, high conversion rates were obtained and the particle size did not significantly affect the reaction conversion. No passivating layer of oxide stopped the oxidation of the core of the particle. The position of Fe2+ in the octa- or tetrahedral site of the cubic mesh has an influence on its reactivity.28 As opposed to FeO, where each iron ion is positioned in an octahedral site, Fe2+ occupies half of the tetrahedral sites in Fe3O4. Thus, Fe2+ is more available for the reaction with a hydroxide in the spinel structure of Fe3O4 than in the cubic mesh of FeO. 4.2. Process Implementation. Experimental results detailed in this paper showed the possibility for the production of hydrogen from three-step cycles based on iron oxides. These cycles are compared in Table 4, with the two-step iron oxide cycle involving the direct hydrolysis of FeO.9 The main drawback of three-step cycles is the corrosiveness of hydroxide compounds, which requires costly construction materials. FeO-based cycles have the solar reduction step in common, which is slow and requires a high temperature and an inert atmosphere. In addition, the required milling of FeO into a fine powder may consume large amounts of energy. The sodium hydroxide reaction may be interesting for fine powders (diameter lower than 20 µm) because of high productivity and energy efficiency. The main advantage of the two-step cycle is the ease of implementation in a process involving only two reactors, solid-gas separations, and no corrosive compound. The three-step cycle based on the Fe2O3/Fe3O4 pair suffers from low hydrogen productivity (4.3 mg of H2/g of Fe3O4) and energy efficiency (25.3%). These drawbacks can be balanced (27) Chen, W. K.; Peterson, N. L. J. Phys. Chem. Solids 1975, 36, 10971103. (28) Gillot, B.; Rousset, A. J. Solid State Chem. 1986, 65, 322-330.

by the simplicity of the reduction of Fe2O3 into Fe3O4, which requires a medium temperature (1300 °C) under air. Thus, a simple nonwindowed solar reactor is sufficient. In addition, reduction into Fe3O4 does not involve a molten phase. Therefore, the particle size remains unchanged during the high-temperature reduction reaction, thereby avoiding a material milling step. The absence of milling and the low reduction temperature may reduce the gap with FeO-based cycles in term of energy efficiency. 5. Summary and Conclusions The three-step Fe2O3/FeO and Fe2O3/Fe3O4 thermochemical cycles achieved hydrogen production, whereas cycles based on cobalt and manganese oxides were not suitable for the cyclic production of H2. The potential to run all reactions involved in iron oxide cycles was demonstrated, and each reactant was regenerated, closing the cycles. Concentrated solar energy suits the temperature requirements for carrying out metal oxide reduction. High-temperature experimental devices were employed to determine reaction temperatures and chemical conversions. Fe2O3 reduction into FeO needed an overheating of liquid Fe3O4 under an inert atmosphere to be complete, and the reaction rate increased when the total pressure was lowered. With the three-step Fe2O3/FeO cycle, the formation of an oxide layer was proven on the FeO surface during the NaOH activation reaction, which stopped hydrogen production. Thus, the chemical conversion was dependent upon the particle size, and a milling of FeO was necessary (particle size below 20 µm) to obtain a chemical conversion higher than 50%. Therefore, the three-step Fe2O3/ FeO cycle was not practical. Next, the three-step cycle involving the Fe2O3/Fe3O4 redox pair was investigated. Fe3O4 was prepared easily with concentrated solar energy under air at 1300 °C. NaOH activation reaction reached a chemical conversion higher than 70% after 7 min at 400 °C, and the final conversion was not dependent upon the particle size in the range of 30-125 µm. The reaction was complete when KOH was employed. Advantages of this cycle include rapid kinetics, high chemical conversion, no milling, and a lower reduction temperature than for FeO (1300 °C under air for Fe3O4 instead of 1600 °C under an inert atmosphere for FeO). Moreover, the hydrolysis of the mixed oxide synthesized in the reaction with NaOH was complete at a temperature of 100 °C, which closed the cycle. Thus, the threestep cycle based on Fe2O3/Fe3O4 appeared to be the most attractive iron-oxide-based cycle for the large-scale production of hydrogen. Acknowledgment. Financial support by the French CNRS (Department of Engineering Science) is gratefully acknowledged. The authors gratefully acknowledge Roger Garcia for construction of the solar experimental reactor and Eric Beche for XRD analysis. EF0701485

Hydrogen Production by Three-Step Solar Thermochemical Cycles

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