Ind. Eng. Chem. Res. 2006, 45, 4993-4998
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Ultrafast Reaction between Li3N and LiNH2 To Prepare the Effective Hydrogen Storage Material Li2NH Yun Hang Hu* and Eli Ruckenstein Department of Chemical Engineering, State UniVersity of New York at Buffalo, Buffalo, New York 14260
Li2NH is a promising candidate for hydrogen storage. However, its performance for hydrogen absorption is strongly dependent on its preparation method. In this paper, we report that an ultrafast solid reaction between Li3N and LiNH2 provides a new synthesis approach to an effective Li2NH material, which can reversibly store 6.8 wt % hydrogen with fast kinetics and excellent stability. In contrast, Li2NH prepared via the conventional LiNH2 decomposition method absorbs less than 2 wt % hydrogen in 500 min. Scanning electron microscopy (SEM) and BET measurements demonstrated that the poor performance of Li2NH prepared via the conventional decomposition method is caused by sintering. In addition, the hydrogen capacity of Li2NH, which was prepared via the reaction between LiH and LiNH2, reached a value of only about 4 wt % after 500 min. 1. Introduction Hydrogen is viewed as a promising clean fuel of the future. For this reason, hydrogen storage technology is critical for the development of a hydrogen-based energy. In recent years, attention has been focused on solid storage materials.1,2 Hydrogen can form hydrides with numerous metals and alloys. However, only some of them can provide reversible hydrogenation-dehydrogenation cycles,3 a requirement essential for hydrogen storage. Although the reaction of hydrides (such as LiH and CaH2) with water has been widely employed to generate hydrogen for meteorological balloons,4 this reaction could not be reversed in any efficient manner. MgH2 exhibits reversible hydrogen storage, but the hydrogenation of magnesium to MgH2 occurs only under severe conditions (high temperatures, above 350 °C, and high pressures, above 50 atm) and very slowly and incompletely. In addition, the rate of dehydrogenation of MgH2 hydride is too low.5 However, the kinetics of hydrogenation could be enhanced through the development of nanocrystalline Mg, Mg alloys, and Mg composites.6-10 High surface area Mg materials (often containing catalysts or other hydriding additives) are now available, which can be rapidly hydrided to >6 wt % hydrogen, even close to room temperature.6 Nevertheless, the dehydrogenation of the hydrogenated nanocrystalline Mg materials still has thermodynamic limitations, which make it hard to recover the H2 at temperatures below 100 °C. Furthermore, it is a great challenge to prevent nanocrystalline Mg from sintering during the hydrogenation-dehydrogenation cycles. The low-temperature reversible hydrides, such as LaNi5H6 and TiFeH2, exhibit suitable dehydrogenation kinetics at very low temperatures;11 they have, however, very low hydrogen storage capacities (1.5 wt % for LaNi5H6 and 1.8 wt % for TiFeH2). The complex hydrides of light metals (Li, Na, and Al), such as LiAlH4 (10.5 wt % H2) and NaAlH4 (7.4 wt % H2), have relatively high hydrogen storage capacities, but they are nonreversible.12 Bogdanovic et al.13-15 demonstrated recently that, by their doping with titanium compounds, the dehydrogenation of NaAlH4/Na3AlH6/Na2LiAlH6 could be facilitated and rendered reversible under moderate conditions. This break* To whom correspondence should be addressed. Tel.: (716) 6452911, ext 2253. Fax: (716) 645-3822. E-mail:
[email protected].
through was followed by progress in the development of catalysts that ensure the reversible dehydrogenation of NaAlH4.16-23 However, a reversible hydrogen storage capacity greater than 6 wt % with good cyclability is still a challenge for these catalyst-doped complex hydrides. Borohydride complexes with suitable alkali or alkaline earth metals also constitute a promising class of compounds for hydrogen storage. Their hydrogen capacity can be as high as 18 wt %. However, to reach a 9 wt % hydrogen desorption, a temperature of 600 °C was required.24-26 It was found that SiO2 mixed with LiBH4 (25:75 by weight) lowered the desorption temperature.27 However, our experiments showed that, although SiO2 can reduce the dehydrogenation temperature of LiBH4, the dehydrogenated LiBH4 could reabsorb only about 1 wt % hydrogen at about 250 °C.28 Other promising materials for hydrogen storage are the nanostructured composites. Hydrogen storage in carbon nanotubes attracted tremendous experimental and theoretical interest.29-37 However, their hydrogen storage is not as efficient as was expected.32-35 For this reason, recent investigations have shifted away from them. Although other types of nanotubes, such as boron nitride,38 MoS2,39 and TiS240 nanotubes, could be used for hydrogen storage, their hydrogen capacities were lower than 3 wt % even at high pressures. Recently, Yaghi and co-workers reported hydrogen storage in microporous metalorganic frameworks.41 Although this observation opened an interesting direction of research, the current microporous metalorganic frameworks have low hydrogen capacities even at 78 K.41b Recently, interest in the hydrogen storage of N-based Li materials has been renewed.42-46 Although the high capacity of hydrogen storage (9.0 wt %) of Li3N could be reproduced,42,47 a critical issue regarding Li3N is that its reversible hydrogen capacity is only about 5.5 wt %.43 This occurs because LiNH2 and 2LiH, which are the products of Li3N hydrogenation, dehydrogenate in two steps,42-44 LiH + LiNH2 ) Li2NH + H2 and LiH + Li2NH ) Li3N + H2. The first step, which provides about 5.5 wt % hydrogen capacity, takes place easily even at temperatures below 200 °C, whereas the second step requires high temperatures (>400 °C). Furthermore, it was found that, although a high dehydrogenation temperature can increase the dehydrogenation, the resulting solid has lost almost completely the ability to reabsorb hydrogen.43b The mole ratio LiNH2/LiH of the hydrogenated Li3N is 0.5, and consequently only half of
10.1021/ie060380i CCC: $33.50 © 2006 American Chemical Society Published on Web 06/08/2006
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LiH releases hydrogen during the first reaction step, at acceptable temperatures. In contrast, if the initial material is lithium imide (Li2NH) instead of Li3N, its hydrogenation products are LiNH2 and LiH (1/1). This means that all LiH in the mixture (LiNH2/LiH) can release hydrogen during the first dehydrogenation step.43,44 As a result, the reversible hydrogen can increase above 6 wt %. For this reason, lithium imide (Li2NH) was considered the most promising because, in principle, it can reversibly absorb 6.85 wt % hydrogen.42,43 The operation temperatures for hydrogen storage are, however, higher than the U.S. Department of Energy target. Nevertheless, it was found that doping Li2NH with Mg or Ca can reduce the dehydrogenation temperature of hydrogenated Li2NH.42,44,46 So far, however, no commercial Li2NH is available. In laboratories, Li2NH is prepared via the direct thermal decomposition of LiNH2 by heating at a temperature of 350 °C (or higher) overnight.42a,45b,c However, this approach requires high energy input because the reaction is endothermic and in addition releases ammonia, which opens environmental issues. Therefore, it is necessary to find an effective approach to prepare Li2NH. Here, we report a fast and effective synthesis approach, in which Li2NH can be generated only in 10 min at 210 °C via the exothermic solid exchange reaction between Li3N and LiNH2 without any byproduct. Furthermore, Li2NH prepared via the fast reaction can reversibly store 6.8 wt % hydrogen with fast kinetics, whereas Li2NH prepared by using the conventional decomposition method absorbs less than 2 wt % hydrogen in 500 min. 2. Material Preparation Design Compared with Li2NH, Li3N has one more Li and one less H, whereas LiNH2 has one more H and one less Li. Therefore, the exchange between the Li of Li3N and the H of LiNH2 is expected to generate Li2NH:
Li3N + LiNH2 ) 2Li2NH
(1)
Furthermore, this reaction is exothermic with ∆H ) -77 kJ/ mol (value that we calculated using the thermodynamic data on Li3N, LiNH2, and Li2NH from ref 42); thus no external energy has to be provided. Starting from this simple observation, we studied the reaction between Li3N and LiNH2 using powders of LiNH2 and Li3N (both bought from Aldrich Chemical Co.) mixed with an agate mortar and pestle by hand for 5 min. The average particle size of the powder, from scanning electron microscopy (Hitachi, S-4000), was about 10 µm. The mixture was subjected to reaction at various temperatures in a vacuum. We denote Li2NH prepared via the above approach as R-Li2NH. For comparison, we also prepared Li2NH via LiNH2 decomposition in a vacuum at various temperatures, which is denoted as β-Li2NH. The reaction LiNH2 + LiH ) Li2NH + H2 was also used to prepare Li2NH (denoted as γ-Li2NH): powders of LiNH2 and LiH (both bought from Aldrich Chemical Co.) were mixed with an agate mortar and pestle by hand for 5 min. The mixture was subjected to reaction at 280 °C in a vacuum overnight. 3.Experimental Section 3.1. Volumetric Test of Hydrogen Storage. The volumetric method, described in a previous paper,43c was employed to accurately determine the hydrogen absorption of Li2NH. A solid sample (0.25 g) was loaded in a reactor located inside an electrical tubular furnace. The change of H2 pressure during
absorption was determined with a digital pressure gauge, which could detect changes in pressure as small as 0.007 atm. The same initial H2 pressure of 7 atm was used in all absorption experiments. An on-line mass spectrometer was used to confirm that, except for hydrogen, no other compounds were present during hydrogenation and dehydrogenation. Before any reabsorption of hydrogen, the sample was subjected to vacuum (p < 10-5 Torr) at 230 °C. It should be noted that the temperature was measured outside the reactor. Therefore, the reaction temperature does not account for the hot spots generated during reaction. The hydrogen capacity is defined as the percentage of hydrogen absorbed based on the total weight of the solid sample (Li2NH). 3.2. BET Surface Areas. A Micromeritics ASAP 2000 instrument was used to determine, via nitrogen adsorption at 77 K, the BET surface areas of various specimens. Because all samples can easily absorb H2O from air, which can increase the surface areas during the degassing process before the BET measurements, we modified the instrument so that all treatments of the samples could be carried out in situ. As a result, we could obtain accurate surface area values. 3.3. X-ray Powder Diffraction (XRD). X-ray powder diffraction of samples were determined using a Siemens D500 X-ray diffraction instrument, equipped with a Cu KR source, at 40 kV and 30 mA. 3.4. Scanning Electron Microscopy (SEM). A scanning electron microscope (Hitachi, S-4000) was employed to examine the morphologies of the specimens. The samples were coated with carbon before measurements. 4. Results and Discussion X-ray powder diffraction measurements for the stoichiometric mixture of Li3N and LiNH2 (1:1 mole ratio) were carried out before and after their reaction. Figure 1a shows that before reaction the mixture contains, as expected, only Li3N and LiNH2. After reaction at 150 °C for 1 h, one can see from Figure 1b that the peaks at 17.4° and 19.6°, which can be attributed to LiNH2, decrease, whereas the peak at about 50°, which can be attributed to Li2NH, increases. The diffraction peaks of LiNH2 and Li2NH between 2θ ) 30° and 50° are very near one another. However, the peaks at 17.4°, 19.6°, and 50° can be used to distinguish LiNH2 from Li2NH. The other five peaks (at 23.1°, 28.4°, 47.2°, 50.4°, and 55.9°), which can be attributed to Li3N, remained present after reaction, indicating that only part of the LiNH2/Li3N was transformed into Li2NH at 150 °C in 1 h. When the reaction took place at 190 °C for 1 h, the LiNH2 phase disappeared and the Li3N phase decreased substantially, whereas the Li2NH phase increased (Figure 1c). When the reaction was carried out at 210 or 230 °C for 1 h, LiNH2 and Li3N were completely transformed into Li2NH (see Figure 1d and Figure 1e, respectively). Surprisingly, at 210 or 230 °C, even 10 min of reaction were enough to completely transform LiNH2 and Li3N into Li2NH (Figure 1f,g), indicating that the reaction was ultrafast at 210 °C or above. This fast reaction can take place by two pathways: (a) gas intermediates and (b) direct ion exchange. The direct ion exchange is unlikely to be dominant, because particles of 10 µm cannot generate large interfaces between them. For this reason, we are inclined to believe that the fast exchange reaction between Li3N and LiNH2 takes place via a gas intermediate: although Li3N decomposition requires a very high temperature (above 813 °C48), LiNH2 can partially decompose to release NH3 even at about 170 °C.42-44 Consequently, at 210 °C or above, the NH3, released from LiNH2,
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Figure 2. Hydrogen absorption by R-Li2NH (prepared via the ultrafast reaction between Li3N and LiNH2) at 230 °C and 7 atm initial hydrogen pressure. Before each reabsorption, the hydrogenated Li2NH was subjected to dehydrogenation at 230 °C for 14 h.
Figure 1. XRD patterns of a stoichiometric mixture of Li3N/LiNH2 (1:1): (a) without any treatment; (b) heated in a vacuum at 150 °C for 1 h; (c) heated in a vacuum at 190 °C for 1 h; (d) heated in a vacuum at 210 °C for 1 h; (e) heated in a vacuum at 230 °C for 1 h; (f) heated in a vacuum at 210 °C for 10 min; (g) heated in a vacuum at 230 °C for 10 min. (Note: Li2O and LiOH were formed because the sample was exposed to air.)
can react with Li3N to form Li2NH:
2LiNH2 ) Li2NH + NH3
(2)
2Li3N + NH3 ) 3Li2NH
(3)
Generally, the direct decomposition of LiNH2 is very slow at temperatures below 350 °C.49,50 This happens because the NH3 equilibrium pressure is very low at temperatures below 350 °C.50 However, in the presence of Li3N, the reaction is much faster, because the capture of NH3 by Li3N (eq 3) reduces the local concentration of NH3, driving the decomposition reaction of LiNH2 to the right (eq 2). Therefore, the reaction between Li3N and LiNH2 in a vacuum at 210-230 °C provides a fast approach to prepare Li2NH, which is denoted as R-Li2NH. The hydrogen absorption by R-Li2NH was determined by using the volumetric method. As shown in Figure 2, the amount of hydrogen absorbed by R-Li2NH reaches 5.4 wt % in 10 min, 6.5 wt % in 60 min, and finally about 6.8 wt %. It is wellknown that Li2NH can easily react with hydrogen at 230 °C to form LiNH2 and LiH (Li2NH + H2 ) LiNH2 + LiH), theoretically absorbing 6.85 wt % hydrogen (based on Li2NH weight).42-44 We also tested the reabsorption of hydrogen after its dehydrogenation at 230 °C for 14 h. It was found that the initial hydrogen capacity first increased with the adsorptiondesorption cycle number and then remained unchanged after four cycles. This means that Li2NH, prepared by the fast reaction between Li3N and LiNH2, is an excellent hydrogen storage material with high capacity and excellent stability. The effect of desorption temperature on hydrogen reabsorption by R-Li2NH was also examined. As shown in Figure 3, the additional desorption at 350 or 450 °C for 3 h after desorption at 230 °C decreases the initial hydrogen capacity. However, the initial hydrogen capacity can be recovered by
Figure 3. Hydrogen reabsorption at 230 °C and 7 atm initial hydrogen pressure by R-Li2NH prepared via the fast reaction between Li3N and LiNH2: (a) first reabsorption after the desorption of hydrogenated Li2NH at 230 °C; (b) first reabsorption after desorption of hydrogenated Li2NH at 230 °C for 14 h and at 350 °C for 3 h; (c) first reabsorption after the desorption of hydrogenated Li2NH at 230 °C for 14 h and at 450 °C for 3 h.
Figure 4. Hydrogen reabsorption at 230 °C and 7 atm initial hydrogen pressure by R-Li2NH. (a) Hydrogen reabsorption after R-Li2NH was subjected to multiple absorption-desorption cycles at 230 °C until hydrogen absorption did not change with the cycle number. (b) Cycl-1, hydrogen reabsorption after R-Li2NH was subjected to vacuum at 230 °C for 14 h and at 450 °C for 3 h; cycl-2, hydrogen reabsorption after the sample used in cycl-1 was subjected to vacuum at 230 °C for 14 h; cycl-3, hydrogen reabsorption after the sample used in cycl-2 was subjected to vacuum at 230 °C for 14 h; cycl-4, hydrogen reabsorption after the sample used in cycl-3 was subjected to vacuum at 230 °C for 14 h.
using several desorption-absorption cycles at 230 °C (Figure 4). Curve a in Figure 4 represents the hydrogen reabsorption by R-Li2NH, which was previously subjected to multiple
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Figure 5. Hydrogen absorption at 230 °C and 7 atm initial hydrogen pressure by β-Li2NH: (a) Li2NH prepared via the conventional decomposition of LiNH2 in a vacuum at 230 °C for overnight; (b) Li2NH prepared via the conventional decomposition of LiNH2 in a vacuum at 280 °C for overnight; (c) Li2NH prepared via the conventional decomposition of LiNH2 in a vacuum at 350 °C for overnight.
absorption-desorption cycles at 230 °C until the hydrogen absorption behavior did not change with cycle number. Curve b cycl-1 represents a much lower initial hydrogen capacity by a R-Li2NH sample, which was previously subjected to vacuum at 230 °C for 14 h and at 450 °C for 3 h, than that of the sample represented by curve a. However, curve b cycl-4 coincides with curve a, indicating that, after four cycles of adsorptiondesorption at 230 °C, the initial hydrogen capacity was completely recovered. For comparison, we also examined the hydrogen absorption by Li2NH (denoted as β-Li2NH), which was prepared by the conventional LiNH2 decomposition method at 230, 280, and 350 °C overnight. One can see from Figure 5 that the reversible hydrogen capacity is less than 2 wt % with slow kinetics. This indicates that the conventional decomposition method for the preparation of Li2NH requires not only heating at high temperatures overnight, which is 100 times longer than that of the fast method, but also produces an ineffective Li2NH for hydrogen storage. In addition, the conventional method releases NH3, which is an air pollutant, and requires high energy input because of its endothermic character. Scanning electron microscopy (SEM) was employed to examine the morphologies of the R- and β-Li2NH samples prepared via the ultrafast reaction between Li3N and LiNH2 and the conventional decomposition of LiNH2 method, respectively. These morphologies are presented in Figures 6 and 7. One can see that the R-Li2NH sample consists of particles, which were about 1 µm (Figure 6). In contrast, the β-Li2NH sample is sintered into blocks (Figure 7). The sintering can explain why β-Li2NH has a much lower hydrogen capacity than R-Li2NH. Furthermore, N2 adsorption was carried out at the temperature of liquid nitrogen (77 K) on the samples. As shown in Table 1, although R-Li2NH has a smaller surface area than its precursor, the Li3N/Li2NH mixture, its surface area remained unchanged when the reaction time was increased from 3 to 9 h at 280 °C. This indicates that the R-Li2NH is stable at 280 °C. However, when the reaction temperature was increased to 350 °C, its surface area decreased from 1.9 to 0.4 m2/g. This can explain why the desorption of the hydrogenated R-Li2NH at higher temperatures led to the reduction of its reabsorption rate for hydrogen. We also measured the surface areas of β-Li2NH prepared via the conventional LiNH2 decomposition method. Table 1 shows that its surface area is only 0.22 and 0.25 m2/g for β-Li2NH prepared at 280 and 350 °C, respectively. This indicates that the material is sintered, which is consistent with
Figure 6. Scanning electron microscopy (SEM) picture for R-Li2NH prepared via the ultrafast reaction between Li3N and LiNH2.
Figure 7. Scanning electron microscopy (SEM) picture for β-Li2NH prepared via the conventional decomposition reaction from LiNH2. Table 1. BET Surface Areas material Li3N/LiNH2 (1/1) R-Li2NH R-Li2NH R-Li2NH R-Li2NH LiNH2 β-Li2NH β-Li2NH
preparation temperature and time
BET surface area (m2/g)
no treatment heating Li3N/LiNH2 (1/1) in a vacuum at 280 °C for 3 h heating Li3N/LiNH2 (1/1) in a vacuum at 280 °C for 6 h heating Li3N/LiNH2 (1/1) in a vacuum at 280 °C for 9 h heating Li3N/LiNH2 (1/1) in a vacuum at 350 °C for 3 h no treatment heating LiNH2 in a vacuum at 280 °C for 3 h heating LiNH2 in a vacuum at 350 °C for 3 h
2.5411 1.9221 1.9576 1.9437 0.397 2.2699 0.2204 0.2548
the SEM measurements. Furthermore, this indicates that the surface area of R-Li2NH is about 10 times larger than that of β-Li2NH, which explains why the hydrogenation of R-Li2NH is much faster than that of β-Li2NH. Another method to prepare Li2NH is the reaction between LiH and LiNH2: LiH + LiNH2 ) Li2NH + H2. As shown in Figure 8, this Li2NH, denoted as γ-Li2NH, prepared via this reaction at 280 °C overnight reaches 4 wt % hydrogen capacity after 500 min of absorption time. Hence, Li2NH prepared via the reaction between LiH and LiNH2 has a much slower kinetics
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Figure 8. Hydrogen absorption at 230 °C and 7 atm initial hydrogen pressure by γ-Li2NH via the reaction of LiH/LiNH2 (1:1) mixture at 280 °C overnight.
than that prepared via the ultrafast reaction between Li3N and LiNH2. Finally, Li2NH can also be prepared by the reaction between NH3 and Li or LiH followed by decomposition. As shown previously, the Li2NH prepared via this reaction has also a low capacity with slow kinetics.43h 4. Conclusions From the above results and discussion, one can conclude that an ultrafast solid reaction between Li3N and LiNH2, which lasts just 10 min, provides a novel synthesis approach to a Li2NH material. Furthermore, this Li2NH material can reversibly store 6.8 wt % hydrogen with fast kinetics and excellent cyclability. In contrast, Li2NH prepared by the conventional LiNH2 decomposition method absorbs less than 2 wt % hydrogen in 500 min. The poor performance of Li2NH prepared via the conventional decomposition method can be attributed to sintering. In addition, the hydrogen capacity of Li2NH, which was prepared via the reaction between LiH and LiNH2, reaches only about 4 wt % after 500 min. Therefore, the ultrafast solid reaction between Li3N and LiNH2 provides an excellent approach to prepare an effective Li2NH material for hydrogen storage. Literature Cited (1) (a) Zuttel, A. Hydrogen Storage materials. Naturwissenschaften 2004, 91, 157, (b) Fichtner, M. Nanotechnological aspects in materials for hydrogen storage. AdV. Eng. Matter 2005, 7, 443. (2) (a) Grochala, W.; Edwards, P. P. Thermal decomposition of the noninterstitial hydrides for the storage and production of hydrogen. Chem. ReV. 2004, 104, 1283. (b) Trudeau, M. L. Advanced materials for energy storage. MRS Bull. 1999, 24, 23. (c) Hu, Y. H.; Ruckenstein, E. A Promising Hydrogen Storage Material;Clathrate Hydrogen Hydrate. Angew. Chem., Int. Ed. 2006, 45, 2011. (3) Sandrock, G. Hydride storage. In Handbook of Fuel Cell; Vielstich, W., Lamm, A., Gasteiger, H. A., Eds.; John Wiley & Sons: New York, 2003; p 101. (4) Gibb, T. R. P., Jr. J. Chem. Educ. 1948, 25, 577. (5) Genossar, J.; Rudman, P. S. Catalytic role of Mg2Cu in the hydriding and dehydriding of Mg. Z. Phys. Chem., N. F. 1979, 116, 215. (6) Huot, J.; Liang, G.; Shultz, R. Mechanically alloyed metal hydride systems. Appl. Phys. A 2001, 72, 187. (7) Orimo, S.; Fujii, H. Materials science of Mg-Ni-based hydrides. Appl. Phys. A 2001, 72, 167. (8) Zaluska, A.; Zaluski, L.; Strom-Olsen, J. O. Structure, catalysis and atomic reactions on the nano-scale: a systematic approach to metal hydrides for hydrogen storage. Appl. Phys. A 2001, 72, 157. (9) Ivanov, E.; Konstanchuk, I.; Bokhonov, B.; Boldyrev, V. Comparative study of first hydriding of Mg-NaF and Mg-NaCl mechanical alloys. J. Alloys Compd. 2003, 360, 256. (10) Tran, N. E.; Imam, M. A.; Feng, C. R. Evaluation of hydrogen storage characteristics of magnesium-misch metal alloys. J. Alloys Compd. 2003, 359, 225. (11) Buchner, H. Energiespeichrung in Metallhydriden; SpringerVerlag: Wien, 1982.
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ReceiVed for reView March 27, 2006 ReVised manuscript receiVed May 5, 2006 Accepted May 9, 2006 IE060380I
