Pd Composites for Hydrogen Separation

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Energy & Fuels 2007, 21, 3530–3536

Stability of Alkali Nitrate/Pd Composites for Hydrogen Separation Membranes J. W. Phair* CSIRO Energy Technology, PriVate Bag 33, Clayton South Victoria, 3169, Australia ReceiVed February 15, 2007. ReVised Manuscript ReceiVed June 12, 2007

A series of alkali nitrate/Pd composites were investigated for their thermal and chemical stability to separate hydrogen using a combination of thermogravimetric and analytical techniques. Alkali nitrates are promising proton conductors which normally do not start to decompose in a reducing environment until temperatures >600 °C have been attained. However, in the presence of Pd, the alkali nitrates undergo decomposition at significantly lower temperatures with decomposition starting at ∼80 °C for NaNO3, 130 °C for KNO3, and between 150 and 200 °C for RbNO3 and CsNO3. The overall decomposition process for the alkali nitrates occurs by reaction with hydrogen, catalyzed by the presence of Pd, which aids the dissociation of H2 f 2H+ + 2e. Decomposition occurs though a series of reactions that reduce the alkali nitrate to NH3, NO2, NO, and alkali metal or alkali metal oxides. In a moist, ambient atmosphere, the alkali metal products (M, M2O) readily convert to their respective metal carbonates. This work demonstrates that screening of promising oxo salts for hydrogen separation membranes by thermogravimetric analysis of the metals in the presence of Pd, the most commonly utilized catalytic layer, is a useful means of assessing their potential stability in hydrogen separation.

Dense ceramic membranes permeable to hydrogen (H2) are potentially useful devices for separating H2 from other gases (e.g., CO, CO2, N2, H2O) formed during energy generation or chemical processing. Hydrogen separation membranes may be applied in the production of hydrogen directly (e.g., the combustion of fossil or biofuels), indirectly (e.g., in integrated coal gasification), or in industrial ecology and process optimization such as during ammonia production. In particular, dense ceramic hydrogen separation membranes are most suited to separate hydrogen under the extreme conditions (500–1100 °C, ∆MPa ∼ 7) encountered during energy generation from coal gasification, where metal, carbon, or polymer membranes are less useful.1,2 Dense ceramic hydrogen separation membranes generally consist of a proton conducting phase, an electron conducting phase, and a catalytic layer for dissociating and reassociating the hydrogen. While the protonic and electronic phases may coexist in a single phase, they are more likely to be present as distinct phases. Initially, dense ceramic membranes based on metal oxide perovskites, for example, BZY (i.e., BaZr0.8Y0.2O3) and BCY (i.e., BaCe0.8Y0.2O3-R) were of greatest interest due to their high proton conductivities at temperatures in excess of 500 °C.3,4 However, the presence of alkali ions as a significant component of the metal oxide increases their susceptibility to react with CO2 or SO2 in the feed gas and form carbonates or sulphates.

This has led to novel innovations in membranes such as proton conducting oxides based on less alkaline rare-earth elements5 or BaCeO3–BaZrO3 solid solutions,6 to eliminate CO2/SO2 susceptibility. Unfortunately, this is usually accompanied by a reduction in proton conductivity. Another approach for developing competitive hydrogen separation membranes is to improve the thermal stability of alternative superprotonic solids (i.e., solids exhibiting high proton conductivity) which are less susceptible to reaction with CO2 or SO2. A major category of these superprotonic conductors are oxy-acid salts which occur naturally in a water-soluble form. To improve their conductivity and mechanical properties, these salts must be stabilized as a ceramic–salt composite and maintained at temperatures greater than 100 °C in order to prevent their dissolution. While the protonic conductivities for many of these materials have been reported,7 less information is available on the critical thermal, chemical, and mechanical properties for application as hydrogen separation membranes. This includes their stability in a hydrogen or syngas environment or interactions with other materials likely to be present in a typical dense hydrogen separation membrane, such as the electron conducting phases or catalytic layers. Alkali nitrate salts have been reported to exhibit high protonic conductivity and can be stabilized at temperatures up to 400 °C, when mixed with a refractory oxide to form a ceramic–salt composite.8–14 One author has reported proton conductivities

* Corresponding author. Tel.: +61 3 9545 2962. Fax: +61 3 9545 2720. E-mail: [email protected]. (1) Guan, J.; Dorris, S. E.; Balachandran, U.; Liu, M. Solid State Ionics 1997, 100, 45–52. (2) Mundschau, M. V.; Xie, X.; Evenson, C. R.; Sammells, A. F. Catal. Today 2006, 118, 12–23. (3) Zuo, C. D.; Lee, T. H.; Dorris, S. E.; Balachandran, U.; Liu, M. L. J. Power Sources 2006, 159, 1291–1295. (4) Zuo, C. D.; Dorris, S. E.; Balachandran, U.; Liu, M. L. Chem. Mater. 2006, 18, 4647–4650.

(5) Haugsrud, R.; Norby, T. Nat. Mater. 2006, 5, 193–196. (6) Ryu, K. H.; Haile, S. M. Solid State Ionics 1999, 125, 355–367. (7) Agrawal, R. C.; Gupta, R. K. J. Mater. Sci. 1999, 34, 1131–1162. (8) Uvarov, N. F.; Vanek, P.; Yuzyuk, Y.; Zelezny, V.; Studnicka, V.; Bokhonov, B. B.; Dulepov, V. E.; Petzelt, J. Solid State Ionics 1996, 90, 201–207. (9) Uvarov, N. F.; Hairetdinov, E. F.; Skobelev, I. V. Solid State Ionics 1996, 86–88, 577–580. (10) Zhu, B.; Mellander, B. E. Solid State Ionics 1994, 70, 285–290. (11) Zhu, B.; Mellander, B. E. J. Power Sources 1994, 52, 289–293.

Introduction

10.1021/ef070086i CCC: $37.00  2007 American Chemical Society Published on Web 09/21/2007

Stability of Alkali Nitrate/Pd Composites

of 0.1–0.01 S cm-1 above 480 °C for RbNO3–Al2O3 composites.15 Another investigation of the proton conduction of RbNO3–SiO2 composites determined conductivities of 0.1 S cm-1 at 280 °C.16 The application of alkali nitrates in hydrogen separation membranes has recently been reported to be advantageous compared to other proton conducting oxy-acid salts (e.g., sulphates and phosphates) due to their ability to function over wider operating temperatures with higher conductivities.17 Nevertheless, doubts remain about their chemical stability in a reducing environment. The purpose of the present work is to report on the stability of alkali nitrates in a hydrogen separation environment mainly based on thermal analysis experiments. Hydrogen separation membranes based on proton conductors typically consist of 30–50 vol % of an electron conducting phase, which is also preferably a hydrogen transport metal for improved efficiency.18 Pd is the most popular and developed transport metal for this role and is also preferred for catalytic layers on membranes.19 Therefore, testing the stability of alkali nitrates in the presence of Pd and in a reducing environment similar to that encountered by hydrogen separation membranes is critical to assessing the potential utility of alkali nitrates in dense hydrogen separation membranes. While previous work has examined the thermal stability of nitrates in ambient or inert environments in the absence of Pd,20–22 no data have been reported on the stability of nitrates in the presence of Pd. Experimental Section The chemical stability of alkali metal nitrates was investigated. The metal nitrates examined were NaNO3, KNO3, RbNO3, and CsNO3. All metal nitrates used were of AR grade or better (purity >99.8%) and obtained from Fluka Chemical Co. Ltd. (Ronkonkoma, NY), Sigma-Aldrich (Milwaukee, WI), or Ajax Chemicals (Sydney NSW, Australia). Sub-micrometer high purity (>99.9+%) Pd powder was obtained from Aldrich Chemical Co. (Milwaukee, WI). Alkali nitrate/Pd composite samples were prepared by crushing, grinding, and then mixing the alkali metal nitrate with 30 wt % Pd with an agate mortar and pestle. After thorough mixing, the samples were dried overnight at 80 °C in a vacuum oven and sealed. To investigate the thermal decomposition of the alkali metal nitrates in a reducing hydrogen separation environment, thermal gravimetric analysis (TGA), differential scanning calorimetry (DSC), and differential thermal analysis (DTA) were conducted under flowing air or H2 utilizing a Setaram SETSYS Evolution TGA-DTA/DSC-1750 or Netzsch TG 409 CD. Ultrahigh purity 10% H2 in a nitrogen gas mixture as received from BOC Gases Australia Ltd. (Ryde NSW, Australia) was used for all experiments. Temperature-dependent measurements (5 °C/min) were performed under a flowing H2 gas rate of 20 cm3/min, and evolved gas analysis (12) Uvarov, N. F.; Skobelev, I. V.; Bokhonov, B. B.; Hairetdinov, E. F. J. Mater. Synth. Process. 2006, 4, 391–395. (13) Rao, M. V. M.; Reddy, S. N.; Chary, A. S.; Shahi, K. Physica B: Condens. Matter 2005, 364, 306–310. (14) Rao, M. V. M.; Reddy, S. N.; Chary, A. S. J. Non-Cryst. Solids 2006, 352, 155–159. (15) Ponomareva, V. G.; Lavrova, G. V.; Simonova, L. G. Solid State Ionics 2000, 136, 1279–1283. (16) Roark, S. E.; Mackay, R.; Mundschau, M. V. Dense, Layered Membranes For Hydrogen Separation. US Patent 7,001,446, 2006. (17) Balachandran, U.; Lee, T. H.; Chen, L.; Song, S. J.; Picciolo, J. J.; Dorris, S. E. Fuel 2006, 85, 150–155. (18) Paglieri, S. N.; Way, J. D. Sep. Purif. Methods 2002, 31, 1–169. (19) Yuvaraj, S.; Fan-Yuan, L.; Tsong-Huei, C.; Chuin-Tih, Y. J. Phys. Chem. B 2003, 107, 1044–1047. (20) Stern, K. H. J. Chem. Educ. 1969, 46, 645. (21) Bond, B. D.; Jacobs, P. W. M. Journal of the Chemical Society A -Inorganic Physical Theoretical 1966, 1265. (22) Liu, H. M.; Hardy, J. R. Phys. ReV. B 1991, 44, 7215–7224.

Energy & Fuels, Vol. 21, No. 6, 2007 3531

Figure 1. Weight change of alkali metal nitrates as a function of temperature under flowing 9% H2 in N2 and a heating rate of 5 °C/ min.

(EGA) was performed using a Pfeiffer Thermostar quadrupole mass spectrometer (MS). Typical sample sizes were 20–40 mg. Prior to each experimental run, the samples were equilibrated at 75 °C under flowing air to eliminate any surface water. The surface water that evolved was monitored by MS, and experiments were conducted only after a steady baseline was achieved. Sample dryness was also confirmed by FTIR measurements. For some of the samples subsequent to the temperature-dependent scans, the chemical stability was further investigated by performing isothermal scans at 200 °C. The H2 flow rate remained at 20 cm3/min. The reaction products of the samples treated in a hydrogen environment were analyzed by X-ray powder diffraction utilizing a Philips 1830/40 X-ray powder diffractometer on finely powdered samples using Cu KR radiation (40 kV and 30 mA) and a Ni filter with a scanning speed of 0.005 °2θ s. The time constant was set at 2 s. To obtain a Fourier transform infrared (FTIR) spectrum, 1.5 mg of sample was ground up with 50 mg of anhydrous KBr and subsequently pressed into a disk. The disk was analyzed in a PerkinElmer 2000 FTIR spectrophotometer, in absorbance mode using KBr as the background reference.

Results and Discussion The thermogravimetry of alkali metal nitrates in air has been well-documented.20–22 Early work by Hogan and Gordon established that increasing the temperature of KNO3,NaNO3, RbNO3, or CsNO3 in air revealed no significant weight loss until temperatures were in excess of 600 °C.23 Data for the thermogravimetry of alkali metal nitrates in the presence of hydrogen are provided in Figure 1. Here, it is clearly apparent that there was no significant weight loss in any of the alkali metal nitrates up to a temperature of 500 °C. This trend is in accordance with a recent confirmation of the decomposition temperatures of NaNO3 and KNO3 in hydrogen as 740 and 840 °C, respectively.19 The reactivity of alkali metal nitrate composites containing 30 wt % Pd in different atmospheres was tested. Samples were treated in air and compared to those treated in a reducing environment (10% H2 in N2). Figure 2 presents the thermogravimetric data obtained for KNO3/Pd treated in both atmospheres up to 400 °C. Clearly, there is a significant difference in thermogravimetric behavior in the two different atmospheres with KNO3/Pd appearing to undergo considerably less weight loss in air than in H2/N2. This suggests that there is no reaction between KNO3 and Pd in air, but in a H2/N2 atmosphere, the KNO3 and Pd appear to react with each other and H2. (23) Hogan, V. D.; Gordon, S. J. Chem. Eng. Data 1961, 6, 572–578.

3532 Energy & Fuels, Vol. 21, No. 6, 2007

Figure 2. Weight change in KNO3 mixed with 30 wt % Pd as a function of temperature under flowing air and 10% H2 in N2 at a heating rate of 5 °C/min.

Figure 3. Heat flow in KNO3 mixed with 30 wt % Pd as a function of temperature under flowing air and 10% H2 in N2 at a heating rate of 5 °C/min.

Figure 3 illustrates the effect of the different atmospheres on the heat flow behavior of the heated KNO3/Pd samples. As for the thermogravimetric data, a significant difference in the heat flow behavior was observed as a function of atmosphere. In air, only two endothermic peaks are observed. The first peak occurring at ∼133 °C is relatively small and correlates with the transition of the room-temperature orthorhombic (space group Pmcn) structure or R-phase of KNO3 to a trigonal structure or β-phase.23 The second and larger endothermic peak occurs at ∼335.7 °C and corresponds to the melting transition of KNO3 from the solid phase. However, in the presence of H2/N2, a new endothermic peak is observed at ∼90 °C which is considerably lower than the first endothermic peak observed in air. Presumably, it is related to a crystal phase transition, but the exact explanation for why it occurs at a lower temperature remains unclear. At ∼127 °C, a slight endothermic peak may be observed but it is quickly overwhelmed by the large exothermic peak between 128 and 180 °C. Most likely, this second endothermic peak is attributed to the same crystal R f β phase transition observed in air. The large exothermic peak is indicative of greater chemical reactivity to be explored shortly. The absence of a melting peak at ∼335 °C in H2/N2 is also interesting given that there had been only ∼10% weight loss by this temperature according to Figure 2. This suggests that there may be phase stabilization occurring in this environment. The thermal stabilities of a series of alkali metal nitrate composites with 30 wt % Pd as a function of temperature in a

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Figure 4. Weight change of alkali metal nitrates mixed with 30 wt % Pd as a function of temperature under flowing 10% H2 in N2 and a heating rate of 5 °C/min.

reducing environment are presented in Figure 4. According to Figure 1, alkali nitrates are relatively stable up to 500 °C in a reducing atmosphere. However, Figure 4 clearly demonstrates the onset of significant weight loss in all of the Pd/alkali metal (Na, K, Rb, Cs) nitrate composites in a reducing environment over a similar temperature range. The decomposition starts at ∼80 °C for NaNO3, ∼130 °C for KNO3, and between 150 and 200 °C for RbNO3 and CsNO3. These temperatures are roughly 600 and 400 °C less than decomposition and reduction temperatures, respectively, reported for alkali nitrates in the absence of palladium.19 The weight loss further increases with increasing temperature. Clearly, the Pd has a distinct effect on the rate and extent of decomposition. In particular, NaNO3 composites demonstrated the greatest weight loss which began at temperatures lower (∼75 °C less) than those observed for the other alkali metal nitrates. The weight loss of NaNO3/Pd is far more rapid than for the other salts, and more than 25% of its original mass has already been lost before the temperature has reached 200 °C. This indicates substantial decomposition of NaNO3 in the reducing environment. While the mass loss for K, Rb, and Cs nitrate is less rapid, decomposition still reaches 10% of the original sample mass within the temperature ranges measured. Explanations for differences in the extent of decomposition of the alkali metal nitrate/Pd composites are not straightforward. While there is a general trend for reducing the extent of decomposition with increasing cation size, further experiments are required to accurately determine the effect of the cation on the mechanism of decomposition. Previously, researchers have compared differences in the charge densities of the cations to infer which compound would have a greater ability to polarize the nitrate ion and therefore form a stronger bond between the nitrate ion and metal.19 However, such a comparison here does not draw a meaningful conclusion. Isothermal scans were performed on the alkali metal nitrate/ Pd samples to confirm the kinetics and extent of weight loss over time. Data for the weight change (%) of KNO3/Pd under a reducing atmosphere over time are provided in Figure 5. In a 10 h period at 200 °C, over 2.5% of the initial mass of the sample was lost with little sign of abatement in the rate of heat loss. In Figure 6, the heat flows of alkali metal nitrate composites with 30 wt % Pd as a function of temperature in a reducing environment are presented. Substantial exothermic peaks are observed for all of the alkali metal nitrate samples indicating the presence of significant chemical reaction. Minor endothermic


Figure 5. Weight loss (percent) isotherm of KNO3 mixed with 30 wt % Pd under flowing 10% H2 in N2 at 200 °C following the initial heating to 400 °C at 5 °C/min then cooling to 200 °C.

Figure 6. Heat flow in alkali metal nitrates mixed with 30 wt % Pd as a function of temperature under flowing 10% H2 in N2 and a heating rate of 5 °C/min.

peaks could also be observed in all samples except NaNO3 in a similar temperature range: 90.7 °C for KNO3, 95.4 °C for RbNO3, 94.6 °C for CsNO3. The exact origin of these endothermic peaks is unclear as they do not correspond to the crystal transition temperatures of the pure alkali nitrates. They could correspond to the stabilization of an alkali nitrate phase which only occurs in the presence of Pd and H2. The major exothermic peaks occur at 144.7 °C for KNO3, 160.8 °C for RbNO3, and 165.3 °C for CsNO3. For NaNO3, however, since the exothermic peak is so broad, it is difficult to determine the exact location of the maximum although the first maximum occurred at 93.8 °C followed by an intermediate maximum at 142.9 °C and a final maximum at 177.2 °C before dropping away steadily. A smaller exothermic peak at 223.6 °C was also observed for the RbNO3 sample. The major trend of these exothermic peaks is that their relative temperature of onset seems to accompany increasing alkali cation size. Interestingly, the main exothermic peaks for all of the alkali metals are significantly closer to the decomposition temperature of the nitrate salt of palladium––Pd(NO3)2 (177 °C)19––than they are to their corresponding nitrate in a reducing environment. At this temperature in a hydrogen-rich environment, Pd2+ is reduced (24) Conner, W. C.; Falconer, J. L. Chem. ReV. 1995, 95, 759–788. (25) Roessner, F.; Roland, U. J. Mol. Catal. A: Chem. 1996, 112, 401– 412.


Figure 7. Heat flow and percent weight change as a function of temperature for CsNO3 mixed with 30 wt % Pd under flowing 10% H2 in N2 and a heating rate of 5 °C/min.

Figure 8. Reproducibility of a typical experiment for the heat flow of RbNO3 mixed with 30 wt % Pd as a function of temperature under flowing 10% H2 in N2 and a heating rate of 5 °C/min.

to Pd and dissociatively absorbs hydrogen which can reduce nitrogen to ammonia through the spillover mechanism.24,25 The main exothermic peak observed during the temperature scan under a reducing environment was accompanied by significant weight loss as demonstrated for CsNO3/Pd in Figure 7. The onset of dramatic weight loss begins just as the major exothermic peak starts to occur at ∼140 °C, suggesting that a chemical reaction is occurring with volatile end products. Weight loss does not subside after the passing of the exothermic peak and continues albeit at a decreasing rate, with increasing temperature. No further major exothermic or endothermic reactions were observed. Multiple runs were performed to confirm experimental reproducibility with multiple runs of RbNO3/Pd presented in Figure 8. The peak positions are in similar positions within experimental error, although a slight shift and variation in intensity of the exothermic peak at ∼220 °C was observed. The reversibility of the reactions was tested by monitoring the heat flow from a sample of RbNO3/Pd in a reducing atmosphere while cooling, as displayed in Figure 9. On heating, the major exothermic peak reached its maximum of 2.6 mW/ mg at 162 °C, with a small shoulder at ∼220 °C. The heat flow was significantly reduced on cooling with the major exothermic peak maximum of 0.20 mW/mg observed at 140 °C. This is followed by a smaller peak at 120 °C with a much reduced heat flow of 0.02 mW/mg at the peak maximum. This confirms (26) Harris, M. J.; Salje, E. K. H.; Guttler, B. K. J. Phys.: Condens. Matter 1990, 2, 5517–5527.


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Figure 9. Heat flow of both the heating and cooling cycles of RbNO3 mixed with 30 wt % Pd as a function of temperature under flowing 10% H2 in N2 and a heating rate of 5 °C/min.

that the reactions are largely irreversible. Furthermore, no endothermic peaks were observed during heating or cooling indicating the absence of any significant phase transitions. A series of chemical reactions that is likely to explain the reduction of alkali metal (M) nitrates in the presence of Pd includes the following: 2MNO3 + 9H2f 98 2M + 2NH3 + 6H2O

(1)

2MNO3 + 5H2f 98 M2O + N2 + 5H2O

(2)

MNO3 + H2f 98 MNO2 + H2O

(3)

2MNO2 + 7H2f 98 2M + 2NH3 + 4H2O

(4)

The main gaseous products of these reactions are H2O, NH3, NO2, and N2. Only, the formation of N2 cannot be directly monitored by the mass spectrometer since the experiments were performed in a N2-rich environment. Temperature-dependent evolved gas analysis during the heating of CsNO3/Pd under a reducing atmosphere is displayed in Figure 10. The ion current is plotted against time for different mass to charge (m/e) integer values and may be construed as the relative concentration of that particular m/e species in the gas derived from the sample. Overall, the mass spectrometry data is in general agreement with the reactions proposed above. However, interpretation of the various m/e channels is complicated by the fact that multiple fragment ions can contribute to the observed intensity for a specific m/e value. For instance, m/e ) 17 can account for both NH3+ and OH+ while m/e ) 16 can consist of contributions from O+ and NH2+, etc. Other likely fragment ions resulting from the reduction of metal nitrates in the presence of Pd include the following: NH+(m/e ) 15), H2O+ (m/e ) 18), N2+ (m/e ) 28), NO+ (m/e ) 30), and NO2+ (m/e ) 46). The exothermic peak maximum for the reduction of CsNO3/ Pd occurs at 165 °C. Concentration peaks occur at 153 °C for NH3+ (m/e ) 16), ∼165 °C for NO2+ (m/e ) 46), and 188 °C for H2O+ (m/e ) 18) as observed in Figure 10a. In addition, a small concentration peak was also observed at ∼165 °C for NO+ (m/e ) 30) although it is not displayed. This confirms that a series of reactions occur upon the heating of CsNO3/Pd in a reducing environment. Special attention needs to be given to the relative concentration of the m/e ) 18 (H2O+) peak which

Figure 10. (a) Mass spectrometry ion current signals for selected m/e channels of interest as a function of temperature for CsNO3 mixed with 30 wt % Pd under flowing 10% H2 in N2 and a heating rate of 5 °C/ min. (b) Mass spectrometry ion current signals for m/e ) 2 and 18 as a function of temperature for CsNO3 mixed with 30 wt % Pd under flowing 10% H2 in N2 and a heating rate of 5 °C/min.

appears to consist of shoulders at ∼165 and 210 °C as well as a broad drop-off in intensity from ∼220 to 370 °C. The shoulder at ∼165 °C may be correlated with the peaks associated with NO2/NO formation while the slow drop-off in intensity may be attributed to ongoing chemical reaction. According to Figure 10b, the evolution of H2O with temperature is accompanied by an almost identical drop in the relative concentration of H2, i.e. hydrogen consumption, with increasing temperature. This indicates that the evolution of H2O is most likely associated with, and a direct result of, the reaction of H2 with CsNO3/Pd rather than any other dominant mechanism. While quantitative analysis of the relative amounts of H2O, NO2/ NO, and NH3 does not allow confirmation of the stoichiometry for any single proposed reaction, it does offer a general confirmation of the species formed. FTIR data were recorded before and after alkali nitrate/Pd composites were treated in a reducing environment (10% H2/ 90% N2) for 12 h at 245 °C. The nitrate anion has a pointgroup symmetry of D3h with the fundamental vibrational modes of A1′, A2′′, and 2E′ of which A2′′ (V2) and 2E′ (V3 and V4) are infrared active. Figure 11 presents the spectra of the nitrates recorded before reduction. For these samples, very little water was detected as indicated by the very small O–H stretching mode at 3200–3500 cm-1, typically associated with water. While



Figure 11. FTIR spectra of Pd/alkali nitrate composites prior to treatment in a reducing atmosphere.

Figure 12. FTIR spectra of Pd/alkali nitrate composites after treatment in a 10% H2 in N2 reducing atmosphere for 12 h at 250 °C.

peak shifts are observed between samples as a function of alkali metal, the major peaks in all samples originate from similar vibrational modes and agree with those reported in the literature for alkali nitrates.26,27 All spectra are dominated by a single sharp peak at ∼1380 cm-1 which can be assigned as V3, the antisymmetric stretch of the nitrate group. The peak at ∼840 cm-1 may be attributed to V2, the out-of-plane bending mode, while the small peak at ∼1800 cm-1 is a combination of the planar deformation (V4) and forbidden symmetric stretching (V1) modes.28 After the alkali nitrate/Pd composites were treated in a reducing environment (10% H2/90% N2) for 12 h at 245 °C, FTIR spectra of the samples were recorded, and the data were (27) Lyndenbell, R. M.; Ferrario, M.; Mcdonald, I. R.; Salje, E. J. Phys.: Condens. Matter 1989, 1, 6523–6542. (28) Brooker, M. H.; Bates, J. B. Spectrochim. Acta, Part A: Molec. Biomolec. Spectrosc. Sect. 1974, 30, 2211–2220. (29) Brooker, M. H.; Bates, J. B. J. Chem. Phys. 1971, 54, 4788–4796.

presented in Figure 12. On initial comparison of the spectra in Figures 11 and Figure 12, there is a significant difference. A strong peak now appears around 700 cm-1; the peak at ∼840 cm-1 shifts to a longer wavenumber as does the peak at ∼1380 cm-1, which also splits into a more complicated structure. The samples also now appear to be considerably wetter owing to the presence of a significant OH stretching peak at 3200 cm-1. The new spectra are clearly different from that of the nitrate anion prior to reduction and may be attributed to the carbonate anion as the predominant species. The unperturbed carbonate anion also has D3h symmetry with A2′′ and 2E′ as the infrared active vibrational modes. At ∼880 cm-1, the out-of-plane deformation V2 (A2′′) is observed while the V3 (E′) and V4(E′) antisymmetric stretch and bend are observed at ∼1400 and 700 cm-1, respectively.29 While the antisymmetric stretch is generally split in the observed spectra, this is not an uncommon feature but may also contain small contributions from unreacted nitrate.


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carbonic acid. The net result is the reaction of carbon dioxide with alkali hydroxide to form an alkali carbonate. Although the experiments were conducted in a dry and reducing environment where the concentration of CO2 would be very low, the samples were exposed to the ambient atmosphere after heat treatment and during the preparation of samples for FTIR/XRD analysis. This would have provided ample opportunity for H2O and CO2 absorption as well as carbonate formation. As a result, one would expect carbonate formation to be highly likely in applications of alkali nitrate salts as membranes for separating H2 from CO2 containing gases. Conclusion

Figure 13. X-ray powder diffraction pattern of the reaction product of NaNO3 with 30 wt % Pd with H2.

Further evidence as to the composition of the alkali nitrate/ Pd samples after reduction could be obtained by X-ray diffraction analysis. Figure 13 provides the X-ray powder diffraction pattern of the product of the reduction of NaNO3 with 30 wt % Pd after a thermogravimetric temperature scan. According to literature values,30 the diffraction peaks match those for sodium carbonate and confirm the results obtained from the FTIR data. As a result of these analytical measurements, it is possible to suggest a mechanism for the formation of the carbonate ions both indirectly and directly which can be represented as follows: M2O + H2O f 2MOH

(5)

2M + 2H2O f 2MOH + H2

(6)

CO2 + H2O f H2CO3

(7a)

2MOH + H2CO3 f M2CO3+2H2O

(7b)

CO2 + 2MOH f M2CO3 + H2O

(8)

Following the reduction of the alkali metal nitrate, amounts of alkali metal or its oxide are left. These are extremely hygroscopic and will react rapidly with any H2O available in the atmosphere to form the metal hydroxide. Upon alkali metal hydroxide formation, which is typically in a wet and concentrated form, it will undergo rapid reaction with carbon dioxide directly or carbon dioxide dissolved from the atmosphere as (30) Dusek, M.; Chapuis, G.; Meyer, M.; Petricek, V. Acta Crystallogr., Sect. B: Struct. Sci. 2003, 59, 337–352.

Alkali nitrates combined with Pd are highly unstable when heated in a reducing environment compared to when in air or in the absence of Pd. The reduction and decomposition of alkali nitrates in a reducing atmosphere occurs at significantly lower temperatures in the presence of Pd, resulting in a complex array of reaction products. Pd catalyzes the reduction of nitrate to ammonia most likely via the spillover mechanism as well as its reduction to NO2/NO. The reduction products of the alkali metal are likely to form carbonates in the presence of CO2. This suggests that the alkali metal nitrates are chemically and thermodynamically unstable when combined with Pd to separate hydrogen in a reducing atmosphere containing CO2. Possible improvements to hydrogen separation membranes based on alkali metal nitrates may rely on the utilization of alternative electron conducting phases and catalytic layers to Pd. However, this may come at the expense of lower hydrogen flux rates. Furthermore, it may be fruitful to utilize different oxo salts, which are considered to be more stable in a reducing environment (e.g., PO43-), as the proton conducting phase in hydrogen separation membranes. A useful test for screening the potential thermal and chemical stability of oxo-anion candidates for hydrogen separation membranes may be to test the stability of the respective Pd salt in a reducing environment and at elevated temperatures. This will provide an indication of the susceptibility of the anion to Pd-catalyzed hydrogen dissociation. Acknowledgment. The author is grateful for the assistance of Mr. P.R. Curtis in the collection of thermogravimetric data, Mrs. L. Goodall for X-ray diffraction analysis, and Dr. J. Mardel for FTIR analysis. Dr. Kate Nairn and Dr. S.P.S Badwal are thanked for reviewing the manuscript. EF070086I

Pd Composites for Hydrogen Separation

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