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Al2O3 for PEM fuel cell applications
Applied Catalysis A: General 227 (2002) 231–240
Ammonia decomposition kinetics over Ni-Pt/Al2 O3 for PEM fuel cell applications A.S. Chellappa1 , C.M. Fischer, W.J. Thomson∗ Department of Chemical Engineering, Washington State University, P.O. Box 642710, Pullman, WA 99164-2710, USA Received 2 July 2001; received in revised form 25 October 2001; accepted 26 October 2001
Abstract Due to the potential employment of ammonia as a hydrogen source for proton exchange membrane (PEM) fuel cell applications, an experimental study has been undertaken to determine the ammonia decomposition rates over a Ni-Pt/Al2 O3 catalyst under conditions likely to be employed in such an application. Using differential analysis of integral data, intrinsic rate data have been obtained between temperatures of 520 and 690 ◦ C and at ammonia pressures between 50 and 780 Torr. It is shown that a first-order rate expression provides an adequate fit of the experimental data over the entire range of pressures and temperatures, although hydrogen inhibition effects may be significant at the lowest temperature. The activation energy was found to be on the order of 50 kcal/mol as opposed to the very low values (5–10 kcal/mol) obtained previously by others. The data are analyzed in the context of previous studies that were conducted under markedly different reaction conditions and differences are attributed to the types of catalysts employed and the fact that previous studies have primarily dealt with very low ammonia concentrations. Consequently, it does not appear as though N–H bond cleavage is the rate determining step under these conditions. It is also shown that the rate expression is capable of predicting, to within 5%, the high ammonia conversions required for fuel cell applications. © 2002 Elsevier Science B.V. All rights reserved. Keywords: Ammonia decomposition; Hydrogen generation; Fuel cells
1. Introduction The drive for reducing emission levels in automobile exhaust and for developing efficient small-scale power generation units (<5 kW) has resulted in a renewed interest in hydrogen generation to power proton exchange membrane (PEM) fuel cells. These fuel cells require a supply of carbon monoxide free hydrogen gas (<50 ppm) to the anode in order to avoid ∗ Corresponding author. Tel.: +1-509-335-8580; fax: +1-509-335-4806. E-mail address: [email protected] (W.J. Thomson). 1 Present address: MesoSystems Technology Inc., 1021 N. Kellogg St., Kennewick, WA 99336, USA.
poisoning of the anode catalyst. Because the cost of hydrogen produced by conventional steam reforming of natural gas is about US$ 5/GJ, and transportation accounts for an additional US$ 10/GJ to the final cost [1], local production of hydrogen (on-site) is considered to be a viable option, particularly if the size of the hydrogen generation unit can be tailored to satisfy the needs of specific power requirements. Currently, the commercialization of fuel cell-based processes has been hampered mainly by two factors, namely, (1) catalyst and reactor stability in an integrated fuel cell system at on-board or on-site hydrogen generation facilities, and (2) delivery of hydrogen generated at industrially sized off-site facilities. The former requires a compact fuel processor, and often
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involves hybrid reactor design, and therefore, has a number of unresolved issues. For instance, while an on-board multi-fuel processor is believed to be a viable option for automobile applications, there are questions related to acceptable transient response characteristics and flexibility with regards to the type of fuel used [2]. Due to the requirement of a virtually carbon monoxide free hydrogen feed, integrated fuel cell processes that use hydrocarbon fuels are rather complex, and resemble a scaled-down refinery unit with multiple processing steps such as desulfurization, water gas shift, methanation, and preferential oxidation to reduce carbon monoxide levels in the reformer outlet gas. Consequently, hydrogen from a single-step source like ammonia decomposition can be an attractive alternative to hydrocarbon fuels for small-scale fuel cell applications. Ammonia can be conveniently stored in liquid form, has a high energy density (3000 Wh/kg), and the safety issues concerning its storage and handling are well established. More importantly, the product stream (hydrogen/nitrogen) is carbon monoxide free. Although ammonia decomposition has been utilized to power alkaline fuel cell prototypes [3,4], concerns about the effect of trace ammonia on the anode, and on the membrane of PEM fuel cells have largely limited its application for PEM fuel cell hydrogen generation. In fact, Uribe et al. [5] have reported that the presence of trace ammonia (as low as 13 ppm) in the hydrogen gas fed to the anode of the PEM cell degrades cell performance. Their results indicated that the platinum catalyst surface (anode) was not directly poisoned by ammonia, but that the decrease in fuel cell performance was caused by the replacement of H+ ions by NH4 + ions within the anode catalyst layer, and also by the decrease in the conductivity of the membrane. While the effect of lower levels of ammonia contamination (ppb range) on PEM cell performance over long periods of continuous operation remains to be determined, it should be noted that the H2 product from both autothermal reforming and partial oxidation of hydrocarbon fuels, generally contains 30–90 ppm ammonia [5]. In any case, these trace amounts of ammonia as well as the equilibrium ammonia concentrations in reactor outlet gases (500–2000 ppm at 600–620 ◦ C, 1 bar) can be reduced to ppb levels by using highly specific adsorbents (Calgon-URC, Grace Davison Grade 514).
Utilizing these concepts, Call et al. [6] have recently demonstrated an integrated 50 W, 1000 Wh ammonia cracker—PEM fuel cell prototype, that contained an adsorber for clean-up of the hydrogen gas. The ammonia concentration in the cracker outlet gas, generally about 2000–3000 ppm, was reduced to <200 ppb by passing the cracker outlet gas through an adsorber. The above facts present a strong argument for the viability of ammonia as a fuel for PEM fuel cell applications. Early ammonia decomposition studies [7–13] were conducted primarily to gain insight into the mechanism of the ammonia synthesis reaction, whereas more recent investigations have addressed the kinetics of the reaction from the standpoint of nitriding processes [14–17], and the abatement of waste streams related to the synthesis of fuel gas from coal [18,19]. In general, these studies were conducted at low ammonia partial pressures (<100 Torr), which was achieved by either using pure ammonia feed at low pressures, or by using a diluted feed (∼10% NH3 ) at higher pressures. Furthermore, when pure ammonia was used, the reaction conditions were so chosen to yield conversions <10%, to satisfy the needs of a differential reactor. Based on these earlier studies, two important conclusions were reached and prompted the motivation behind this study: 1. At low temperatures and high hydrogen partial pressures, the reaction is inhibited by hydrogen, and is described by the Temkin–Pyzhev mechanism [13], while at high temperatures and low hydrogen partial pressure, the reaction is only dependent on ammonia partial pressure. This behavior has been referred to as the Tamaru model [15], after Tamaru [13], who first made this distinction based on the findings of several investigations. 2. At low temperatures (<500 ◦ C for Pt and Ru) and low ammonia partial pressure (typically <1 Torr) the reaction is zero-order with respect to ammonia, and at high temperatures the reaction becomes first-order with respect to ammonia. This transition was clearly observed by Löffler and Schmidt [9,10], while studying the decomposition of ammonia over platinum and iron. For fuel cell applications, pure ammonia needs to be almost completely decomposed (>99.5% conversion) at around atmospheric pressure. To achieve this degree of conversion, high temperatures are required
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for both thermodynamic and kinetic reasons, and the quantitative rate expressions must apply over a wide range of ammonia partial pressures where hydrogen inhibition effects are likely to be encountered. Since previous studies have not addressed the kinetics of ammonia decomposition under these conditions (most of the studies utilized low ammonia concentrations), this study was undertaken to find out if the mechanistic schemes and the kinetic parameters found in the literature can be utilized when pure ammonia is decomposed to produce hydrogen for fuel cells. For this purpose, the reaction temperature was systematically varied from 520 to 660 ◦ C, with selected experiments also run at temperatures as high as 690 ◦ C, and while using pure ammonia as feed at total pressures up to 2 bar. The reaction rates were then examined in the context of proposed rate expressions in order to arrive at a quantitative rate model appropriate for predictive purposes.
2. Experimental 2.1. Equipment and experimental procedures The decomposition of ammonia was studied in an integral flow reactor fed with pure ammonia and utilized a quartz reactor (10 mm i.d.) attached to a RXM-100 Catalyst Testing and Characterization System (Advanced Scientific Design). Ammonia flow rates were in the 10–500 sccm range, which produced catalyst loading-to-ammonia flow rates (W/F) between
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0.3 and 57 (g cat. h)/g mol NH3 . The absolute pressure measured at the reactor inlet varied from 0.96 to 2.0 bar and the catalyst-bed temperatures ranged from 520 to 690 ◦ C. Anhydrous ammonia (Air Products, 99.99% min), hydrogen (Air Liquide, 99.99% min), and helium (Air Liquide, Pre-purified) were metered through Brooks 5850E mass flow controllers. The effluent gases were analyzed by gas chromatography (Carle Analytical) using a 6 ft × 1/8 in. Porapak N column (Alltech Associates, 80/100 mesh) and a thermal conductivity detector. Before venting, the outlet gas mixture was scrubbed using a 1 M sulfuric acid solution. The platinum oxide-nickel oxide/Al2 O3 catalyst (G43-A, United Catalyst) was crushed and sized to 40 m particles. This catalyst was chosen based on preliminary screening experiments (Table 1), and was found to be relatively more active than the other catalysts tested. Pre-treatment of the catalysts consisted of heating in helium (75 sccm) at 400 ◦ C for 30 min, reducing with hydrogen (55 sccm) at 800 ◦ C for 60 min, and cooling to the desired reaction temperature under hydrogen flow. Dynamic high-temperature X-ray diffraction (DXRD) was used to identify the required reduction temperature. These in situ reduction experiments were performed in a Philips X’PERT-MPD system, which was equipped with a PW3050/10 θ –θ goniometer, a Co K␣1 (1.7890 Å) source of radiation, a flowthrough HTK 1200 Anton Paar oven and a Raytech position-sensitive detector. DXRD scans of the freshly reduced and post-reaction G43-A catalyst samples showed no evidence of nitride formation
Table 1 Catalyst screening experiments conducted at W/F = 3.4a Catalystb
Composition
T (◦ C)
XNH3 c (%)
G43-A (United Catalyst)
1–5 wt.% nickel oxide, <1 wt.% platinum oxide in alumina
500 600
14.5 78.1
2800 (Grace Davison)
Raney nickel 93.8 wt.% Ni, 0.3 wt.% Fe, 5.9 wt.% Al
600 700
18.8 81.6
146 (Johnson Matthey)
0.5% Ru in alumina
500 550 600 700
6.5 12.9 23.6 84.5
a
(g cat. h)/g mol NH3 . 125 m particles. c Equilibrium conversion at 500 ◦ C, 1 bar = 99.74%. b
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Table 2 a Parameters for the r = k0 exp( − Ea /RT) × pNH model for ammonia decomposition 3 Catalyst
T (◦ C)a
Tr (◦ C)a
pNH3 (Torr)
Ea (kcal/mol)
a
Reference
Polycrystalline Fe wire Polycrystalline Pt wire Ru (0 0 1) crystal Polycrystalline Pt wire Nickel wire Vanadium nitride TiNx Oy Fe/Fe4 N Vanadium carbides Ni-Pt/Al2 O3
327–977 227–1427 227–977 127–927 477–1077 446–863 607–666 325–500 380–570 520–660
<677, <571, <477, <280, <727, <620, NA NA NA NA
0.05–1 0.015–10.4 1 × 10−6 to 2 × 10−6 0.5–2 × 10−6 0.04–1.0 159–602 4–22 152–621 760 100–780
49.6, 10 21, 4.3 43, 6.6 22, 4 50.4, 5.9 33, 4.4 33.2 21 30–38 46.9
0, 1 0, 1 0, 1 0, 1 0, 1 0, 1 0.75 1 1 1
[9] [10] [11] [12] [25] [14] [15] [16] [26] This work
a
>677 >571 >477 >280 >727 >620
T and Tr represent the reaction temperature and transition temperature (zero- to first-order kinetics), respectively.
under the reaction conditions employed in this study. The BET surface area (measured using a Coulter SA 3100 apparatus) of the freshly reduced and post-reaction G43-A catalyst was 158 m2 /g and 125 ± 2 m2 /g, respectively. This reduction in surface area is probably an effect of sintering, but did not result in any decrease in catalyst activity with reaction time (as will be shown later). The GC technique provided quantitative data for all three components—N2 , H2 and NH3 . Thus, it was possible to close the material balance for each experiment, and the results were found to be accurate to within 10%. The raw data were used to calculate mole fractions for each component (yi ), and ammonia conversion (XNH3 ), was calculated as XNH3 =
2yN2 2yN2 + yNH3
(1)
yielding plots of conversion versus W/F at each temperature. These integral conversion data were then fit with polynomials to generate equations that could be analytically differentiated to obtain ammonia decomposition rates as a function of conversion (and hence species partial pressures). Because of flow rate limitations, it was not possible to do differential analysis of the data above 660 ◦ C. Consequently, these data were analyzed by comparing them to the predictions of the rate expressions obtained at the lower temperatures. When rate expressions are reported, the reaction rate, r, is in g mol NH3 /(g cat. h), the partial pressures are in bar, and the apparent activation energy, Ea , is in cal/mol.
2.2. Kinetic analyses As mentioned earlier, ammonia decomposition kinetics are available in the literature, but are based on experiments that were conducted at low ammonia concentrations and/or low conversions—conditions that were far different from that required for fuel cell applications. Those studies, which are summarized in Table 2, primarily used two separate rate models [13,15,20,21] to describe ammonia decomposition kinetics. A brief summary of these rate expressions is presented below to provide a comparative basis for the kinetic analysis conducted in this study. At low ammonia concentrations and/or low conversions, ammonia decomposition in these studies was effectively described, using the following sequence: 1. (a) activation of ammonia k1
NH3 + 2∗ → ∗ NH2 + ∗ H
(2)
2. (b) recombinative N2 desorption k2
2∗ N → N2 + 2∗
(3)
where symbol (∗) is a free active site, and ∗ N, ∗ NH2 the surface adsorbed species. Step (a) proceeds to (b) through a number of kinetically insignificant steps, and the two steps are assumed to be far from equilibrium. Assuming adsorbed ammonia (∗ N) to be the most abundant reactive intermediate, that is, L ∼ = ∗ + ∗ N, where L is the total site density, and that nitrogen
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desorption is rate limiting, the ammonia decomposition rate is given by Eq. (4): r=
k1 LpNH3 (1 + (k1 /k2 pNH3 )1/2 )2
(4)
Previous work has shown that, at low temperatures, the decomposition rate was independent of ammonia pressure, suggesting that k2 k1 and r ∝ k2 . Not surprisingly, the low temperature region was characterized by apparent activation energies, which are consistent with the activation energies reported for the recombinative desorption of nitrogen (30–50 kcal/mol). On the other hand, at high temperatures the decomposition rate was observed to be first-order in ammonia partial pressure (r ∝ k1 pNH3 ), with activation energies in the 4–10 kcal/mol range. These values are very close to the calculated energies for N–H bond cleavage. As shown in Table 2, such temperature dependencies have been demonstrated [9–12,14] at low ammonia pressures (<100 Torr) and in the 300–1400 ◦ C range. Ammonia conversions in these studies were typically <10%, and the decomposition rate was found to be independent of nitrogen and hydrogen partial pressures. Furthermore, as pointed out by Oyama [14] and by Djèga–Mariadassou et al. [15], a rate expression that is equivalent to Eq. (4) can be derived by applying Langmuir–Hinshelwood analysis to a sequence that involves equilibrium adsorption of ammonia (Eq. (5)), followed by decomposition of adsorbed ammonia to products (Eq. (6)). K
NH3 (g) ↔ NH3 (s),
ka K= kd
kR
NH3 (s) → 21 N2 (g) + 23 H2 (g)
(5) (6)
yielding, r=
kR k pNH3 , 1 + k pNH3
k =
ka kd + k R
(7)
In this case, the activation energies in the limiting cases also have values similar to those mentioned earlier, as listed in Table 2 [9–12]. A fundamental difference is that, with this expression, the high temperature (first-order) region has apparent activation energies that are a combination of the energies of dissociation and adsorption/desorption of ammonia. Oyama [14] also attempted to account for surface non-uniformity and derived a two-step rate expression
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as a function of the width of the non-uniform surface distribution. This analysis led to a rate expression that was one-half-order with respect to ammonia and a rate constant that was inversely proportional to the width. However, he observed considerable deviation between his calculated (non-uniform rate expression) and experimental results over a vanadium nitride catalyst, and concluded that such a rate expression was not appropriate for decomposition kinetics. At high hydrogen partial pressures, and particularly at low temperatures (<300 ◦ C) other studies [10,13,15,22] have found hydrogen to have an inhibitive effect on the decomposition rate when hydrogen was co-fed along with the ammonia, as could occur during tail gas waste treatment. This behavior has been explained by employing a reaction sequence that involves the quasi-equilibrated step shown in Eq. (8), followed by a rate-limiting nitrogen desorption step (Eq. (3)): K
2[NH3 +∗ ⇔3 ∗ N + 23 H2 ]
(8)
With the same assumptions used to derive Eq. (4), the rate of decomposition is given by r=
2 k2 LK23 pNH 3 3/2
[K3 pNH3 + pH2 ]2
(9)
All of these rate expressions can also be formulated in terms of a power law rate model; i.e. a r = k pNH pb 3 H2
(10)
where k , a and b are constants. For example, at high temperatures (low ammonia partial pressures), both Eqs. (4) and (7) become first-order with respect to ammonia and zero-order with respect to hydrogen. Similarly, when hydrogen inhibition is significant, (high hydrogen partial pressures and low temperatures), b in Eq. (10) is negative, and Eq. (10) becomes equivalent to the Temkin–Pyzhev expression: 2 β pNH3
r=k (11) 3 pH 2 where, k
and β are constants, with β = 0.25–0.6 for iron catalysts of various compositions [23]. Note that all these expressions ignore the reverse reaction, which is appropriate for the experimental conditions employed, which are far removed from equilibrium
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(equilibrium ammonia conversion at 500 ◦ C and 1 bar is 99.74%).
3. Results and discussion To insure repeatability, 2–3 separate GC samples were taken and averaged for each experimental data point, and analyses were typically within ±5% of each other. In addition, in order to determine whether the experiments were repeatable as a whole, repeat runs were also conducted, using different catalyst loadings. Fig. 1 shows these results and, as can be seen, repeatability was excellent over a wide range of flow rates and for three separate catalyst loadings. In addition, catalyst stability over extended periods of time, was assessed by periodically returning to identical experimental conditions whenever the catalyst was idled (over night, weekends, etc.). Fig. 2 shows the results of one such set of experiments, and it can be seen that the catalyst remained stable over a 15-day time period and over a range of flow rates, indicating that the catalyst does not lose activity. These results imply that the effect of sintering on catalyst activity is negligible. Furthermore, to eliminate bulk and interparticle diffusion effects on the reaction rate, separate experiments were carried out with different catalyst particle sizes, and also with different catalyst loadings and flow rates (the latter to give the same space velocity). The results
Fig. 1. Repeatability with three different catalyst loadings at 520 ◦ C. Loading 1 (䉫), loading 2 (䊉) and loading 3 (䊐).
Fig. 2. Catalyst stability measured at 540 ◦ C and using loading #2 at the start of each experiment over a 15-day period. Day 1 (䉬), day 9 (䊐), day 10 (), day 14 (×), day 15 (䊊).
of these experiments that were conducted at mass flow rates of 0.8–14 g/(cm2 h) are shown in Fig. 3. As can be seen, there is no effect of either particle size or flow rate for particle sizes <125 m, and even at the lower end of the range of flow rates employed in the study. Intraparticle diffusion resistances were also estimated, according to the Weisz–Prater criterion [24],
Fig. 3. Effect of particle size and linear velocity on conversion at 600 ◦ C. The catalyst loadings and particle size were: 0.32 g, 40 m (), 0.64 g, 40 m (䊊), 0.15 g, 125 m (䉬), 0.32 g, 125 m (䉱).
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Fig. 4. Representative W/F vs. ammonia conversion plots: 560 ◦ C ( ), 580 ◦ C (䊉), 620 ◦ C (䉱) and 660 ◦ C (䉬).
and were found to be negligible at the highest reaction rate measured in this study. Because of the high temperatures used, and the high ammonia conversions obtained in this study, the experimental data (71 data points measured between 520 and 660 ◦ C) were first analyzed by performing multiple linear regression on the generic power law model as shown in Eq. (12), with the range of ammonia and hydrogen pressures being 100–780 Torr and 150–600 Torr, respectively. −Ea a (12) r = k0 exp pNH pb 3 H2 RT For this purpose, the reaction rates were calculated by differential analysis of the integral data, and representative results of the data are shown in Fig. 4. The statistical analysis showed that the only significant partial pressure variable was that of ammonia, indicating that hydrogen had no effect on the reaction rate. The resulting power law rate expression is −46897 (13) pNH3 r = 3.639 × 1011 exp RT with the rate being given in g mol NH3 /(g cat. h) and the partial pressures are given in bar. Fig. 5 shows a parity plot for this expression and, as can be seen, an excellent fit is obtained, with the only major deviation being at the lowest temperature of 520 ◦ C. The first-order dependency on ammonia is inconsistent with the predicted effects of surface non-uniformity
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Fig. 5. Parity plot of experimental and calculated reaction rates (corresponding to Eq. (13)): 520 ◦ C (䉬), 540 ◦ C ( ), 560 ◦ C (䉱), 580 ◦ C (䊊), 600 ◦ C (䊉), 620 ◦ C (+), 640 ◦ C (), 660 ◦ C (䉫).
[14] and, thus, it appears that the reaction rate does not depend on structural features such as crystallographic plane or particle size of the active catalyst component. Oyama [14] also observed considerable deviation between his calculated (non-uniform rate expression) and experimental results over vanadium nitride catalyst, and concluded that such a rate expression was not appropriate for decomposition kinetics. It is possible that hydrogen inhibition is playing a role at the lower temperatures, but without co-feeding hydrogen with the ammonia, the data are too highly correlated to be able to statistically distinguish a separate effect of hydrogen. Nevertheless, it should be noted that when the data between 520 and 560 ◦ C were analyzed separately with the Temkin–Pyzhev model, Eq. (11), the reaction orders of ammonia and hydrogen were 0.67 and −1, respectively. So there is a definite possibility that hydrogen inhibition could be significant <520 ◦ C, even though these temperatures are higher than the temperatures at which this phenomenon has been typically observed [8,21]. However, Tsai et al. [25] found that ammonia pressure had a significant effect on the temperature where hydrogen inhibition became important and, thus, the higher ammonia pressures used in this work may very well explain the possible occurrence of hydrogen inhibition at 520 ◦ C.
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As can be seen from Table 2, the apparent activation energy of 47 kcal/mol for this Pt-Ni/Al2 O3 catalyst is much higher than most of the previous results which were obtained when the ammonia reaction order was first-order. In fact, high activation energies are normally associated in the zero-order reaction order range [9,11,14,26]. However, with the exception of Oyama [14] all of these investigations were conducted at very low ammonia pressures. On the other hand, Oyama employed a vanadium nitride catalyst and it may be that this phenomenon is catalyst specific. Some evidence for this hypothesis can be found by noting that others have also reported higher activation energies in the first-order reaction regime over other nitrides and carbides [15,16,27]. Only McCabe [26] employed a nickel catalyst (wires), but his ammonia pressures were very low (0.04–1.0 Torr). While Choi [27] is the only author who specifically mentions that mass transfer effects were absent, Oyama’s published flow rates appear to be comparable to ours, although his catalyst was different. Note also, that the very low activation energies were always associated with a clear transition between zero order and first-order, although the temperatures of the transition varied over a wide range (200 ◦ C). None of the previous authors who report high activation energies in the first-order region, observed a transition in the reaction order. In the context of the previous mechanistic schemes proposed by others [13–15,20,21], it could be hypothesized that the absence of a changing reaction order in our case is that ammonia activation over the Ni-Pt/Al2 O3 catalyst follows the quasi-equilibrated path (Eq. (5)) as opposed to Eq. (2). While this
normally results in hydrogen inhibition (Temkin– Pyzhev kinetics), adsorption coverage by hydrogen would be expected to be much less at the higher temperatures employed in our study. A more likely possibility is that ammonia decomposition is governed by more than one rate-limiting step. For example, Bradford et al. [21] point out that the rate expression shown in Eq. (9), cannot simultaneously account for experimentally observed near-first-order ammonia dependence and hydrogen inhibitive effects, and they found that a rate expression involving both NH2 –H bond cleavage and recombinative nitrogen desorption, provided a good fit to their low temperature data over a Ru/C catalyst. If ammonia is to be used as the hydrogen source for fuel cells, both thermodynamics and kinetics dictates that the fuel processor be operated at high temperatures. Because equipment limitations precluded differential analysis of the data above 660 ◦ C, the rate expression given by Eq. (13) was integrated to predict the ammonia conversions under all experimental conditions where the conversion was ≥80%, including data at temperatures up to 690 ◦ C. The integrated form of the rate equation is simply, 1 − XNH3 = exp( − kf (W/F )) where kf is given by kf = 3.639 × 1011 exp
(14)
−46897 RT
(15)
The comparison of these predictions with the experimental measurements is shown in Table 3. As can be seen, the high conversions predicted by Eq. (13) were
Table 3 Comparison of measured and calculated ammonia conversionsa T (◦ C)
W/F (g cat. h)/g molNH3
X (measured)
Xb (calculated)
Xc (calculated)
540 560 580 600 620 640 660 670 690
56.89, 28.44, 8.13 4.89 1.96, 0.78, 0.49 0.58, 0.33
0.974, 0.990, 0.876 0.821 0.898, 0.823, 0.906 0.918, 0.961
0.994, 0.994, 0.943 0.961 0.907, 0.813, 0.842 0.941, 0.933
0.987, 0.989, 0.926 0.952 0.899, 0.812, 0.848 0.947, 0.943
40.63 18.97
3.26 0.98, 1.96, 1.39 0.49
X ≥ 0.80. Differential analysis (520–660 ◦ C) data. c Integral analysis (520–690 ◦ C) data. a
b
0.864 0.916
0.966 0.834, 0.995, 0.928 0.896
0.974 0.967
0.981 0.879, 0.986, 0.951 0.910
0.956 0.951
0.978 0.878, 0.985, 0.951 0.918
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largely within 10% of the experimental values and the extrapolation to temperatures >660 ◦ C was also reasonable. Because integration is an inherent smoothing process, it was decided to submit all the data to an integral analysis, assuming first-order with respect to ammonia. When this was done, the rate constant was found to be −49229 kf = 1.309 × 1012 exp (16) RT A comparison of the high conversion predictions of Eq. (16) shows that the predicted values are somewhat improved, with all but one data point within 5% of the experimental values. The good correspondence between the experimental and predicted values clearly shows that the simple first-order model is quite robust in satisfactorily predicting conversions >80% over a wide range of W/F values and operating temperatures (W/F = 0.3 to 57 g cat. h/g mol NH3 , T = 540–690 ◦ C). This study has therefore demonstrated that first-order decomposition kinetics can be reliably integrated into the design of a ammonia cracker for fuel cells. Also, in view of the fact that the activation energy for this Ni-Pt/Al2 O3 catalyst is closer to that reported for nickel, it appears that platinum behaves more like a stabilizer in preventing the sintering of nickel active sites, rather than as a catalyst promoter.
4. Conclusions This investigation of ammonia decomposition for hydrogen production has demonstrated that the kinetics are quite different than those reported previously at low ammonia concentrations. Under high ammonia concentrations and high temperatures, a rate expression that is first-order with respect to ammonia appears to be adequate for purposes of predicting ammonia conversions under conditions corresponding to fuel processors for PEM fuel cells. In contrast to previous work, first-order behavior with respect to ammonia was observed over the entire temperature and ammonia pressure range. It is concluded that the transition from zero-order to first-order kinetics does not apply under these conditions, and it is probable that hydrogen inhibition is influenced by both ammonia pressure and temperature. In addition, the
239
activation energy measured here is much higher (ca. 50 versus 5–10 kcal/mol) than that reported by earlier workers operating at high temperatures and low ammonia pressures. A comparison of these results with other values reported in the literature indicates that the activation energy appears to be catalyst specific. Mechanistically, this would imply that N–H bond cleavage is not the rate determining step under these conditions. Finally, it is concluded that a simple first-order rate expression, with the rate constant given by Eq. (16), is applicable to ammonia decomposition at temperatures >520 ◦ C and pressures above 100 Torr, when using a Ni-based catalyst. At these conditions, there is no hydrogen inhibition and there is no apparent change in the reaction order with respect to ammonia. It should be pointed out that extrapolation of the first-order rate expression should be done with caution; particularly to lower temperatures and ammonia pressures. In addition, the prediction of ammonia conversions in excess of 0.99 would require an accounting of the reversible nature of this reaction.
Acknowledgements InnovaTek Inc. is gratefully acknowledged for partial financial support as well as for the efforts of Dr. Lloyd Allen, who assisted in the initial catalyst screening. We also wish to thank United Catalyst Inc. for supplying the catalyst, K. Taylor for helping with some of the preliminary experiments, and M.R. Powell for discussions related to the effects of trace ammonia on PEM fuel cell performance. References [1] R. Socolow (Ed.), Fuels Decarbonization and Carbon Sequestration, Report No. PU/CEES 302, Princeton University Press, Princeton, September 1997, p. 17. [2] P.S. Chintawar, C. Papile, W.L. Mitchell, in: Proceedings of the 16th Meeting of the North American on Catalyst Society, Boston, MA, June 1999. [3] K. Kordesch, V. Hacker, J. Gsellmann, M. Cifrain, G. Faleschini, P. Enzinger, M. Ortner, M. Muhr, R.R. Aronson, in: Proceedings of the 3rd International Fuel Cell Conference, Nagoya, Japan, November 1999. [4] R. Metkemeijer, P. Achard, J. Power Sources 49 (1994) 271. [5] F.A. Uribe, T. Zawodzinski Jr, S. Gottesfeld, Electrochem. Soc. Meet. Abstr. 98 (2) (1998) 1103.
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[6] C.J. Call, M.R. Powell, M. Fountain, A.S. Chellappa, in: Proceedings of the Knowledge Foundation’s 3rd Annual International Symposium on Fuel Cells for Portable Power Applications, Washington, DC, April 2001. [7] G.C. Bond, Catalysis by Metals, Academic Press, London, 1962, p. 371. [8] A. Amano, H. Taylor, J. Am. Chem. Soc. 76 (1954) 4201. [9] D.G. Löffler, L.D. Schmidt, J. Catal. 44 (1976) 244. [10] D.G. Löffler, L.D. Schmidt, J. Catal. 41 (1976) 440. [11] W. Tsai, W.H. Weinberg, J. Phys. Chem. 91 (1987) 5302. [12] J.J. Vajo, W. Tsai, W.H. Weinberg, J. Phys. Chem. 89 (1985) 3243. [13] K. Tamaru, Acc. Chem. Res. 21 (1988) 88. [14] S.T. Oyama, J. Catal. 133 (1992) 358. [15] G. Djèga-Mariadassou, C.-H. Shin, G. Bugli, J. Mol. Catal. A: Chem. 141 (1999) 263. [16] W. Arabczyk, J. Zamlynny, Catal. Lett. 60 (1999) 167. [17] R.S. Wise, E.J. Markel, J. Catal. 145 (1994) 335.
[18] K. Jothimurugesan, S.K. Gangwal, in: Proceedings of the US DOE Advanced Coal-Fired Power Systems’96 Review Meeting, Morgantown, WV, July 1996. [19] W. Wang, N. Padban, Z. Ye, A. Andersson, I. Bjerle, Ind. Eng. Chem. Res. 38 (1999) 4175. [20] M. Boudart, G. Djèga-Mariadassou, Kinetics of Heterogeneous Catalytic Reactions, Princeton University Press, Princeton, NJ, 1984, p. 98. [21] M.C.J. Bradford, P.E. Fanning, M.A. Vannice, J. Catal. 172 (1997) 479. [22] G. Ertl, M. Huber, J. Catal. 61 (1980) 537. [23] C.N. Satterfield, Heterogeneous Catalysis in Industrial Practice, McGraw-Hill, New York, 1991, p. 464. [24] P.B. Weisz, C.D. Prater, Adv. Catal. 6 (1954) 143. [25] W. Tsai, J.J. Vajo, W.H. Weinberg, J. Phys. Chem. 89 (1985) 4926. [26] R.W. McCabe, J. Catal. 79 (1983) 445. [27] J.-G. Choi, J. Catal. 182 (1999) 104.
Ammonia decomposition kinetics over Ni-Pt/Al2 O3 for PEM fuel cell applications A.S. Chellappa1 , C.M. Fischer, W.J. Thomson∗ Department of Chemical Engineering, Washington State University, P.O. Box 642710, Pullman, WA 99164-2710, USA Received 2 July 2001; received in revised form 25 October 2001; accepted 26 October 2001
Abstract Due to the potential employment of ammonia as a hydrogen source for proton exchange membrane (PEM) fuel cell applications, an experimental study has been undertaken to determine the ammonia decomposition rates over a Ni-Pt/Al2 O3 catalyst under conditions likely to be employed in such an application. Using differential analysis of integral data, intrinsic rate data have been obtained between temperatures of 520 and 690 ◦ C and at ammonia pressures between 50 and 780 Torr. It is shown that a first-order rate expression provides an adequate fit of the experimental data over the entire range of pressures and temperatures, although hydrogen inhibition effects may be significant at the lowest temperature. The activation energy was found to be on the order of 50 kcal/mol as opposed to the very low values (5–10 kcal/mol) obtained previously by others. The data are analyzed in the context of previous studies that were conducted under markedly different reaction conditions and differences are attributed to the types of catalysts employed and the fact that previous studies have primarily dealt with very low ammonia concentrations. Consequently, it does not appear as though N–H bond cleavage is the rate determining step under these conditions. It is also shown that the rate expression is capable of predicting, to within 5%, the high ammonia conversions required for fuel cell applications. © 2002 Elsevier Science B.V. All rights reserved. Keywords: Ammonia decomposition; Hydrogen generation; Fuel cells
1. Introduction The drive for reducing emission levels in automobile exhaust and for developing efficient small-scale power generation units (<5 kW) has resulted in a renewed interest in hydrogen generation to power proton exchange membrane (PEM) fuel cells. These fuel cells require a supply of carbon monoxide free hydrogen gas (<50 ppm) to the anode in order to avoid ∗ Corresponding author. Tel.: +1-509-335-8580; fax: +1-509-335-4806. E-mail address: [email protected] (W.J. Thomson). 1 Present address: MesoSystems Technology Inc., 1021 N. Kellogg St., Kennewick, WA 99336, USA.
poisoning of the anode catalyst. Because the cost of hydrogen produced by conventional steam reforming of natural gas is about US$ 5/GJ, and transportation accounts for an additional US$ 10/GJ to the final cost [1], local production of hydrogen (on-site) is considered to be a viable option, particularly if the size of the hydrogen generation unit can be tailored to satisfy the needs of specific power requirements. Currently, the commercialization of fuel cell-based processes has been hampered mainly by two factors, namely, (1) catalyst and reactor stability in an integrated fuel cell system at on-board or on-site hydrogen generation facilities, and (2) delivery of hydrogen generated at industrially sized off-site facilities. The former requires a compact fuel processor, and often
0926-860X/02/$ – see front matter © 2002 Elsevier Science B.V. All rights reserved. PII: S 0 9 2 6 - 8 6 0 X ( 0 1 ) 0 0 9 4 1 - 3
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involves hybrid reactor design, and therefore, has a number of unresolved issues. For instance, while an on-board multi-fuel processor is believed to be a viable option for automobile applications, there are questions related to acceptable transient response characteristics and flexibility with regards to the type of fuel used [2]. Due to the requirement of a virtually carbon monoxide free hydrogen feed, integrated fuel cell processes that use hydrocarbon fuels are rather complex, and resemble a scaled-down refinery unit with multiple processing steps such as desulfurization, water gas shift, methanation, and preferential oxidation to reduce carbon monoxide levels in the reformer outlet gas. Consequently, hydrogen from a single-step source like ammonia decomposition can be an attractive alternative to hydrocarbon fuels for small-scale fuel cell applications. Ammonia can be conveniently stored in liquid form, has a high energy density (3000 Wh/kg), and the safety issues concerning its storage and handling are well established. More importantly, the product stream (hydrogen/nitrogen) is carbon monoxide free. Although ammonia decomposition has been utilized to power alkaline fuel cell prototypes [3,4], concerns about the effect of trace ammonia on the anode, and on the membrane of PEM fuel cells have largely limited its application for PEM fuel cell hydrogen generation. In fact, Uribe et al. [5] have reported that the presence of trace ammonia (as low as 13 ppm) in the hydrogen gas fed to the anode of the PEM cell degrades cell performance. Their results indicated that the platinum catalyst surface (anode) was not directly poisoned by ammonia, but that the decrease in fuel cell performance was caused by the replacement of H+ ions by NH4 + ions within the anode catalyst layer, and also by the decrease in the conductivity of the membrane. While the effect of lower levels of ammonia contamination (ppb range) on PEM cell performance over long periods of continuous operation remains to be determined, it should be noted that the H2 product from both autothermal reforming and partial oxidation of hydrocarbon fuels, generally contains 30–90 ppm ammonia [5]. In any case, these trace amounts of ammonia as well as the equilibrium ammonia concentrations in reactor outlet gases (500–2000 ppm at 600–620 ◦ C, 1 bar) can be reduced to ppb levels by using highly specific adsorbents (Calgon-URC, Grace Davison Grade 514).
Utilizing these concepts, Call et al. [6] have recently demonstrated an integrated 50 W, 1000 Wh ammonia cracker—PEM fuel cell prototype, that contained an adsorber for clean-up of the hydrogen gas. The ammonia concentration in the cracker outlet gas, generally about 2000–3000 ppm, was reduced to <200 ppb by passing the cracker outlet gas through an adsorber. The above facts present a strong argument for the viability of ammonia as a fuel for PEM fuel cell applications. Early ammonia decomposition studies [7–13] were conducted primarily to gain insight into the mechanism of the ammonia synthesis reaction, whereas more recent investigations have addressed the kinetics of the reaction from the standpoint of nitriding processes [14–17], and the abatement of waste streams related to the synthesis of fuel gas from coal [18,19]. In general, these studies were conducted at low ammonia partial pressures (<100 Torr), which was achieved by either using pure ammonia feed at low pressures, or by using a diluted feed (∼10% NH3 ) at higher pressures. Furthermore, when pure ammonia was used, the reaction conditions were so chosen to yield conversions <10%, to satisfy the needs of a differential reactor. Based on these earlier studies, two important conclusions were reached and prompted the motivation behind this study: 1. At low temperatures and high hydrogen partial pressures, the reaction is inhibited by hydrogen, and is described by the Temkin–Pyzhev mechanism [13], while at high temperatures and low hydrogen partial pressure, the reaction is only dependent on ammonia partial pressure. This behavior has been referred to as the Tamaru model [15], after Tamaru [13], who first made this distinction based on the findings of several investigations. 2. At low temperatures (<500 ◦ C for Pt and Ru) and low ammonia partial pressure (typically <1 Torr) the reaction is zero-order with respect to ammonia, and at high temperatures the reaction becomes first-order with respect to ammonia. This transition was clearly observed by Löffler and Schmidt [9,10], while studying the decomposition of ammonia over platinum and iron. For fuel cell applications, pure ammonia needs to be almost completely decomposed (>99.5% conversion) at around atmospheric pressure. To achieve this degree of conversion, high temperatures are required
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for both thermodynamic and kinetic reasons, and the quantitative rate expressions must apply over a wide range of ammonia partial pressures where hydrogen inhibition effects are likely to be encountered. Since previous studies have not addressed the kinetics of ammonia decomposition under these conditions (most of the studies utilized low ammonia concentrations), this study was undertaken to find out if the mechanistic schemes and the kinetic parameters found in the literature can be utilized when pure ammonia is decomposed to produce hydrogen for fuel cells. For this purpose, the reaction temperature was systematically varied from 520 to 660 ◦ C, with selected experiments also run at temperatures as high as 690 ◦ C, and while using pure ammonia as feed at total pressures up to 2 bar. The reaction rates were then examined in the context of proposed rate expressions in order to arrive at a quantitative rate model appropriate for predictive purposes.
2. Experimental 2.1. Equipment and experimental procedures The decomposition of ammonia was studied in an integral flow reactor fed with pure ammonia and utilized a quartz reactor (10 mm i.d.) attached to a RXM-100 Catalyst Testing and Characterization System (Advanced Scientific Design). Ammonia flow rates were in the 10–500 sccm range, which produced catalyst loading-to-ammonia flow rates (W/F) between
233
0.3 and 57 (g cat. h)/g mol NH3 . The absolute pressure measured at the reactor inlet varied from 0.96 to 2.0 bar and the catalyst-bed temperatures ranged from 520 to 690 ◦ C. Anhydrous ammonia (Air Products, 99.99% min), hydrogen (Air Liquide, 99.99% min), and helium (Air Liquide, Pre-purified) were metered through Brooks 5850E mass flow controllers. The effluent gases were analyzed by gas chromatography (Carle Analytical) using a 6 ft × 1/8 in. Porapak N column (Alltech Associates, 80/100 mesh) and a thermal conductivity detector. Before venting, the outlet gas mixture was scrubbed using a 1 M sulfuric acid solution. The platinum oxide-nickel oxide/Al2 O3 catalyst (G43-A, United Catalyst) was crushed and sized to 40 m particles. This catalyst was chosen based on preliminary screening experiments (Table 1), and was found to be relatively more active than the other catalysts tested. Pre-treatment of the catalysts consisted of heating in helium (75 sccm) at 400 ◦ C for 30 min, reducing with hydrogen (55 sccm) at 800 ◦ C for 60 min, and cooling to the desired reaction temperature under hydrogen flow. Dynamic high-temperature X-ray diffraction (DXRD) was used to identify the required reduction temperature. These in situ reduction experiments were performed in a Philips X’PERT-MPD system, which was equipped with a PW3050/10 θ –θ goniometer, a Co K␣1 (1.7890 Å) source of radiation, a flowthrough HTK 1200 Anton Paar oven and a Raytech position-sensitive detector. DXRD scans of the freshly reduced and post-reaction G43-A catalyst samples showed no evidence of nitride formation
Table 1 Catalyst screening experiments conducted at W/F = 3.4a Catalystb
Composition
T (◦ C)
XNH3 c (%)
G43-A (United Catalyst)
1–5 wt.% nickel oxide, <1 wt.% platinum oxide in alumina
500 600
14.5 78.1
2800 (Grace Davison)
Raney nickel 93.8 wt.% Ni, 0.3 wt.% Fe, 5.9 wt.% Al
600 700
18.8 81.6
146 (Johnson Matthey)
0.5% Ru in alumina
500 550 600 700
6.5 12.9 23.6 84.5
a
(g cat. h)/g mol NH3 . 125 m particles. c Equilibrium conversion at 500 ◦ C, 1 bar = 99.74%. b
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Table 2 a Parameters for the r = k0 exp( − Ea /RT) × pNH model for ammonia decomposition 3 Catalyst
T (◦ C)a
Tr (◦ C)a
pNH3 (Torr)
Ea (kcal/mol)
a
Reference
Polycrystalline Fe wire Polycrystalline Pt wire Ru (0 0 1) crystal Polycrystalline Pt wire Nickel wire Vanadium nitride TiNx Oy Fe/Fe4 N Vanadium carbides Ni-Pt/Al2 O3
327–977 227–1427 227–977 127–927 477–1077 446–863 607–666 325–500 380–570 520–660
<677, <571, <477, <280, <727, <620, NA NA NA NA
0.05–1 0.015–10.4 1 × 10−6 to 2 × 10−6 0.5–2 × 10−6 0.04–1.0 159–602 4–22 152–621 760 100–780
49.6, 10 21, 4.3 43, 6.6 22, 4 50.4, 5.9 33, 4.4 33.2 21 30–38 46.9
0, 1 0, 1 0, 1 0, 1 0, 1 0, 1 0.75 1 1 1
[9] [10] [11] [12] [25] [14] [15] [16] [26] This work
a
>677 >571 >477 >280 >727 >620
T and Tr represent the reaction temperature and transition temperature (zero- to first-order kinetics), respectively.
under the reaction conditions employed in this study. The BET surface area (measured using a Coulter SA 3100 apparatus) of the freshly reduced and post-reaction G43-A catalyst was 158 m2 /g and 125 ± 2 m2 /g, respectively. This reduction in surface area is probably an effect of sintering, but did not result in any decrease in catalyst activity with reaction time (as will be shown later). The GC technique provided quantitative data for all three components—N2 , H2 and NH3 . Thus, it was possible to close the material balance for each experiment, and the results were found to be accurate to within 10%. The raw data were used to calculate mole fractions for each component (yi ), and ammonia conversion (XNH3 ), was calculated as XNH3 =
2yN2 2yN2 + yNH3
(1)
yielding plots of conversion versus W/F at each temperature. These integral conversion data were then fit with polynomials to generate equations that could be analytically differentiated to obtain ammonia decomposition rates as a function of conversion (and hence species partial pressures). Because of flow rate limitations, it was not possible to do differential analysis of the data above 660 ◦ C. Consequently, these data were analyzed by comparing them to the predictions of the rate expressions obtained at the lower temperatures. When rate expressions are reported, the reaction rate, r, is in g mol NH3 /(g cat. h), the partial pressures are in bar, and the apparent activation energy, Ea , is in cal/mol.
2.2. Kinetic analyses As mentioned earlier, ammonia decomposition kinetics are available in the literature, but are based on experiments that were conducted at low ammonia concentrations and/or low conversions—conditions that were far different from that required for fuel cell applications. Those studies, which are summarized in Table 2, primarily used two separate rate models [13,15,20,21] to describe ammonia decomposition kinetics. A brief summary of these rate expressions is presented below to provide a comparative basis for the kinetic analysis conducted in this study. At low ammonia concentrations and/or low conversions, ammonia decomposition in these studies was effectively described, using the following sequence: 1. (a) activation of ammonia k1
NH3 + 2∗ → ∗ NH2 + ∗ H
(2)
2. (b) recombinative N2 desorption k2
2∗ N → N2 + 2∗
(3)
where symbol (∗) is a free active site, and ∗ N, ∗ NH2 the surface adsorbed species. Step (a) proceeds to (b) through a number of kinetically insignificant steps, and the two steps are assumed to be far from equilibrium. Assuming adsorbed ammonia (∗ N) to be the most abundant reactive intermediate, that is, L ∼ = ∗ + ∗ N, where L is the total site density, and that nitrogen
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desorption is rate limiting, the ammonia decomposition rate is given by Eq. (4): r=
k1 LpNH3 (1 + (k1 /k2 pNH3 )1/2 )2
(4)
Previous work has shown that, at low temperatures, the decomposition rate was independent of ammonia pressure, suggesting that k2 k1 and r ∝ k2 . Not surprisingly, the low temperature region was characterized by apparent activation energies, which are consistent with the activation energies reported for the recombinative desorption of nitrogen (30–50 kcal/mol). On the other hand, at high temperatures the decomposition rate was observed to be first-order in ammonia partial pressure (r ∝ k1 pNH3 ), with activation energies in the 4–10 kcal/mol range. These values are very close to the calculated energies for N–H bond cleavage. As shown in Table 2, such temperature dependencies have been demonstrated [9–12,14] at low ammonia pressures (<100 Torr) and in the 300–1400 ◦ C range. Ammonia conversions in these studies were typically <10%, and the decomposition rate was found to be independent of nitrogen and hydrogen partial pressures. Furthermore, as pointed out by Oyama [14] and by Djèga–Mariadassou et al. [15], a rate expression that is equivalent to Eq. (4) can be derived by applying Langmuir–Hinshelwood analysis to a sequence that involves equilibrium adsorption of ammonia (Eq. (5)), followed by decomposition of adsorbed ammonia to products (Eq. (6)). K
NH3 (g) ↔ NH3 (s),
ka K= kd
kR
NH3 (s) → 21 N2 (g) + 23 H2 (g)
(5) (6)
yielding, r=
kR k pNH3 , 1 + k pNH3
k =
ka kd + k R
(7)
In this case, the activation energies in the limiting cases also have values similar to those mentioned earlier, as listed in Table 2 [9–12]. A fundamental difference is that, with this expression, the high temperature (first-order) region has apparent activation energies that are a combination of the energies of dissociation and adsorption/desorption of ammonia. Oyama [14] also attempted to account for surface non-uniformity and derived a two-step rate expression
235
as a function of the width of the non-uniform surface distribution. This analysis led to a rate expression that was one-half-order with respect to ammonia and a rate constant that was inversely proportional to the width. However, he observed considerable deviation between his calculated (non-uniform rate expression) and experimental results over a vanadium nitride catalyst, and concluded that such a rate expression was not appropriate for decomposition kinetics. At high hydrogen partial pressures, and particularly at low temperatures (<300 ◦ C) other studies [10,13,15,22] have found hydrogen to have an inhibitive effect on the decomposition rate when hydrogen was co-fed along with the ammonia, as could occur during tail gas waste treatment. This behavior has been explained by employing a reaction sequence that involves the quasi-equilibrated step shown in Eq. (8), followed by a rate-limiting nitrogen desorption step (Eq. (3)): K
2[NH3 +∗ ⇔3 ∗ N + 23 H2 ]
(8)
With the same assumptions used to derive Eq. (4), the rate of decomposition is given by r=
2 k2 LK23 pNH 3 3/2
[K3 pNH3 + pH2 ]2
(9)
All of these rate expressions can also be formulated in terms of a power law rate model; i.e. a r = k pNH pb 3 H2
(10)
where k , a and b are constants. For example, at high temperatures (low ammonia partial pressures), both Eqs. (4) and (7) become first-order with respect to ammonia and zero-order with respect to hydrogen. Similarly, when hydrogen inhibition is significant, (high hydrogen partial pressures and low temperatures), b in Eq. (10) is negative, and Eq. (10) becomes equivalent to the Temkin–Pyzhev expression: 2 β pNH3
r=k (11) 3 pH 2 where, k
and β are constants, with β = 0.25–0.6 for iron catalysts of various compositions [23]. Note that all these expressions ignore the reverse reaction, which is appropriate for the experimental conditions employed, which are far removed from equilibrium
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(equilibrium ammonia conversion at 500 ◦ C and 1 bar is 99.74%).
3. Results and discussion To insure repeatability, 2–3 separate GC samples were taken and averaged for each experimental data point, and analyses were typically within ±5% of each other. In addition, in order to determine whether the experiments were repeatable as a whole, repeat runs were also conducted, using different catalyst loadings. Fig. 1 shows these results and, as can be seen, repeatability was excellent over a wide range of flow rates and for three separate catalyst loadings. In addition, catalyst stability over extended periods of time, was assessed by periodically returning to identical experimental conditions whenever the catalyst was idled (over night, weekends, etc.). Fig. 2 shows the results of one such set of experiments, and it can be seen that the catalyst remained stable over a 15-day time period and over a range of flow rates, indicating that the catalyst does not lose activity. These results imply that the effect of sintering on catalyst activity is negligible. Furthermore, to eliminate bulk and interparticle diffusion effects on the reaction rate, separate experiments were carried out with different catalyst particle sizes, and also with different catalyst loadings and flow rates (the latter to give the same space velocity). The results
Fig. 1. Repeatability with three different catalyst loadings at 520 ◦ C. Loading 1 (䉫), loading 2 (䊉) and loading 3 (䊐).
Fig. 2. Catalyst stability measured at 540 ◦ C and using loading #2 at the start of each experiment over a 15-day period. Day 1 (䉬), day 9 (䊐), day 10 (), day 14 (×), day 15 (䊊).
of these experiments that were conducted at mass flow rates of 0.8–14 g/(cm2 h) are shown in Fig. 3. As can be seen, there is no effect of either particle size or flow rate for particle sizes <125 m, and even at the lower end of the range of flow rates employed in the study. Intraparticle diffusion resistances were also estimated, according to the Weisz–Prater criterion [24],
Fig. 3. Effect of particle size and linear velocity on conversion at 600 ◦ C. The catalyst loadings and particle size were: 0.32 g, 40 m (), 0.64 g, 40 m (䊊), 0.15 g, 125 m (䉬), 0.32 g, 125 m (䉱).
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Fig. 4. Representative W/F vs. ammonia conversion plots: 560 ◦ C ( ), 580 ◦ C (䊉), 620 ◦ C (䉱) and 660 ◦ C (䉬).
and were found to be negligible at the highest reaction rate measured in this study. Because of the high temperatures used, and the high ammonia conversions obtained in this study, the experimental data (71 data points measured between 520 and 660 ◦ C) were first analyzed by performing multiple linear regression on the generic power law model as shown in Eq. (12), with the range of ammonia and hydrogen pressures being 100–780 Torr and 150–600 Torr, respectively. −Ea a (12) r = k0 exp pNH pb 3 H2 RT For this purpose, the reaction rates were calculated by differential analysis of the integral data, and representative results of the data are shown in Fig. 4. The statistical analysis showed that the only significant partial pressure variable was that of ammonia, indicating that hydrogen had no effect on the reaction rate. The resulting power law rate expression is −46897 (13) pNH3 r = 3.639 × 1011 exp RT with the rate being given in g mol NH3 /(g cat. h) and the partial pressures are given in bar. Fig. 5 shows a parity plot for this expression and, as can be seen, an excellent fit is obtained, with the only major deviation being at the lowest temperature of 520 ◦ C. The first-order dependency on ammonia is inconsistent with the predicted effects of surface non-uniformity
237
Fig. 5. Parity plot of experimental and calculated reaction rates (corresponding to Eq. (13)): 520 ◦ C (䉬), 540 ◦ C ( ), 560 ◦ C (䉱), 580 ◦ C (䊊), 600 ◦ C (䊉), 620 ◦ C (+), 640 ◦ C (), 660 ◦ C (䉫).
[14] and, thus, it appears that the reaction rate does not depend on structural features such as crystallographic plane or particle size of the active catalyst component. Oyama [14] also observed considerable deviation between his calculated (non-uniform rate expression) and experimental results over vanadium nitride catalyst, and concluded that such a rate expression was not appropriate for decomposition kinetics. It is possible that hydrogen inhibition is playing a role at the lower temperatures, but without co-feeding hydrogen with the ammonia, the data are too highly correlated to be able to statistically distinguish a separate effect of hydrogen. Nevertheless, it should be noted that when the data between 520 and 560 ◦ C were analyzed separately with the Temkin–Pyzhev model, Eq. (11), the reaction orders of ammonia and hydrogen were 0.67 and −1, respectively. So there is a definite possibility that hydrogen inhibition could be significant <520 ◦ C, even though these temperatures are higher than the temperatures at which this phenomenon has been typically observed [8,21]. However, Tsai et al. [25] found that ammonia pressure had a significant effect on the temperature where hydrogen inhibition became important and, thus, the higher ammonia pressures used in this work may very well explain the possible occurrence of hydrogen inhibition at 520 ◦ C.
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As can be seen from Table 2, the apparent activation energy of 47 kcal/mol for this Pt-Ni/Al2 O3 catalyst is much higher than most of the previous results which were obtained when the ammonia reaction order was first-order. In fact, high activation energies are normally associated in the zero-order reaction order range [9,11,14,26]. However, with the exception of Oyama [14] all of these investigations were conducted at very low ammonia pressures. On the other hand, Oyama employed a vanadium nitride catalyst and it may be that this phenomenon is catalyst specific. Some evidence for this hypothesis can be found by noting that others have also reported higher activation energies in the first-order reaction regime over other nitrides and carbides [15,16,27]. Only McCabe [26] employed a nickel catalyst (wires), but his ammonia pressures were very low (0.04–1.0 Torr). While Choi [27] is the only author who specifically mentions that mass transfer effects were absent, Oyama’s published flow rates appear to be comparable to ours, although his catalyst was different. Note also, that the very low activation energies were always associated with a clear transition between zero order and first-order, although the temperatures of the transition varied over a wide range (200 ◦ C). None of the previous authors who report high activation energies in the first-order region, observed a transition in the reaction order. In the context of the previous mechanistic schemes proposed by others [13–15,20,21], it could be hypothesized that the absence of a changing reaction order in our case is that ammonia activation over the Ni-Pt/Al2 O3 catalyst follows the quasi-equilibrated path (Eq. (5)) as opposed to Eq. (2). While this
normally results in hydrogen inhibition (Temkin– Pyzhev kinetics), adsorption coverage by hydrogen would be expected to be much less at the higher temperatures employed in our study. A more likely possibility is that ammonia decomposition is governed by more than one rate-limiting step. For example, Bradford et al. [21] point out that the rate expression shown in Eq. (9), cannot simultaneously account for experimentally observed near-first-order ammonia dependence and hydrogen inhibitive effects, and they found that a rate expression involving both NH2 –H bond cleavage and recombinative nitrogen desorption, provided a good fit to their low temperature data over a Ru/C catalyst. If ammonia is to be used as the hydrogen source for fuel cells, both thermodynamics and kinetics dictates that the fuel processor be operated at high temperatures. Because equipment limitations precluded differential analysis of the data above 660 ◦ C, the rate expression given by Eq. (13) was integrated to predict the ammonia conversions under all experimental conditions where the conversion was ≥80%, including data at temperatures up to 690 ◦ C. The integrated form of the rate equation is simply, 1 − XNH3 = exp( − kf (W/F )) where kf is given by kf = 3.639 × 1011 exp
(14)
−46897 RT
(15)
The comparison of these predictions with the experimental measurements is shown in Table 3. As can be seen, the high conversions predicted by Eq. (13) were
Table 3 Comparison of measured and calculated ammonia conversionsa T (◦ C)
W/F (g cat. h)/g molNH3
X (measured)
Xb (calculated)
Xc (calculated)
540 560 580 600 620 640 660 670 690
56.89, 28.44, 8.13 4.89 1.96, 0.78, 0.49 0.58, 0.33
0.974, 0.990, 0.876 0.821 0.898, 0.823, 0.906 0.918, 0.961
0.994, 0.994, 0.943 0.961 0.907, 0.813, 0.842 0.941, 0.933
0.987, 0.989, 0.926 0.952 0.899, 0.812, 0.848 0.947, 0.943
40.63 18.97
3.26 0.98, 1.96, 1.39 0.49
X ≥ 0.80. Differential analysis (520–660 ◦ C) data. c Integral analysis (520–690 ◦ C) data. a
b
0.864 0.916
0.966 0.834, 0.995, 0.928 0.896
0.974 0.967
0.981 0.879, 0.986, 0.951 0.910
0.956 0.951
0.978 0.878, 0.985, 0.951 0.918
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largely within 10% of the experimental values and the extrapolation to temperatures >660 ◦ C was also reasonable. Because integration is an inherent smoothing process, it was decided to submit all the data to an integral analysis, assuming first-order with respect to ammonia. When this was done, the rate constant was found to be −49229 kf = 1.309 × 1012 exp (16) RT A comparison of the high conversion predictions of Eq. (16) shows that the predicted values are somewhat improved, with all but one data point within 5% of the experimental values. The good correspondence between the experimental and predicted values clearly shows that the simple first-order model is quite robust in satisfactorily predicting conversions >80% over a wide range of W/F values and operating temperatures (W/F = 0.3 to 57 g cat. h/g mol NH3 , T = 540–690 ◦ C). This study has therefore demonstrated that first-order decomposition kinetics can be reliably integrated into the design of a ammonia cracker for fuel cells. Also, in view of the fact that the activation energy for this Ni-Pt/Al2 O3 catalyst is closer to that reported for nickel, it appears that platinum behaves more like a stabilizer in preventing the sintering of nickel active sites, rather than as a catalyst promoter.
4. Conclusions This investigation of ammonia decomposition for hydrogen production has demonstrated that the kinetics are quite different than those reported previously at low ammonia concentrations. Under high ammonia concentrations and high temperatures, a rate expression that is first-order with respect to ammonia appears to be adequate for purposes of predicting ammonia conversions under conditions corresponding to fuel processors for PEM fuel cells. In contrast to previous work, first-order behavior with respect to ammonia was observed over the entire temperature and ammonia pressure range. It is concluded that the transition from zero-order to first-order kinetics does not apply under these conditions, and it is probable that hydrogen inhibition is influenced by both ammonia pressure and temperature. In addition, the
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activation energy measured here is much higher (ca. 50 versus 5–10 kcal/mol) than that reported by earlier workers operating at high temperatures and low ammonia pressures. A comparison of these results with other values reported in the literature indicates that the activation energy appears to be catalyst specific. Mechanistically, this would imply that N–H bond cleavage is not the rate determining step under these conditions. Finally, it is concluded that a simple first-order rate expression, with the rate constant given by Eq. (16), is applicable to ammonia decomposition at temperatures >520 ◦ C and pressures above 100 Torr, when using a Ni-based catalyst. At these conditions, there is no hydrogen inhibition and there is no apparent change in the reaction order with respect to ammonia. It should be pointed out that extrapolation of the first-order rate expression should be done with caution; particularly to lower temperatures and ammonia pressures. In addition, the prediction of ammonia conversions in excess of 0.99 would require an accounting of the reversible nature of this reaction.
Acknowledgements InnovaTek Inc. is gratefully acknowledged for partial financial support as well as for the efforts of Dr. Lloyd Allen, who assisted in the initial catalyst screening. We also wish to thank United Catalyst Inc. for supplying the catalyst, K. Taylor for helping with some of the preliminary experiments, and M.R. Powell for discussions related to the effects of trace ammonia on PEM fuel cell performance. References [1] R. Socolow (Ed.), Fuels Decarbonization and Carbon Sequestration, Report No. PU/CEES 302, Princeton University Press, Princeton, September 1997, p. 17. [2] P.S. Chintawar, C. Papile, W.L. Mitchell, in: Proceedings of the 16th Meeting of the North American on Catalyst Society, Boston, MA, June 1999. [3] K. Kordesch, V. Hacker, J. Gsellmann, M. Cifrain, G. Faleschini, P. Enzinger, M. Ortner, M. Muhr, R.R. Aronson, in: Proceedings of the 3rd International Fuel Cell Conference, Nagoya, Japan, November 1999. [4] R. Metkemeijer, P. Achard, J. Power Sources 49 (1994) 271. [5] F.A. Uribe, T. Zawodzinski Jr, S. Gottesfeld, Electrochem. Soc. Meet. Abstr. 98 (2) (1998) 1103.
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[6] C.J. Call, M.R. Powell, M. Fountain, A.S. Chellappa, in: Proceedings of the Knowledge Foundation’s 3rd Annual International Symposium on Fuel Cells for Portable Power Applications, Washington, DC, April 2001. [7] G.C. Bond, Catalysis by Metals, Academic Press, London, 1962, p. 371. [8] A. Amano, H. Taylor, J. Am. Chem. Soc. 76 (1954) 4201. [9] D.G. Löffler, L.D. Schmidt, J. Catal. 44 (1976) 244. [10] D.G. Löffler, L.D. Schmidt, J. Catal. 41 (1976) 440. [11] W. Tsai, W.H. Weinberg, J. Phys. Chem. 91 (1987) 5302. [12] J.J. Vajo, W. Tsai, W.H. Weinberg, J. Phys. Chem. 89 (1985) 3243. [13] K. Tamaru, Acc. Chem. Res. 21 (1988) 88. [14] S.T. Oyama, J. Catal. 133 (1992) 358. [15] G. Djèga-Mariadassou, C.-H. Shin, G. Bugli, J. Mol. Catal. A: Chem. 141 (1999) 263. [16] W. Arabczyk, J. Zamlynny, Catal. Lett. 60 (1999) 167. [17] R.S. Wise, E.J. Markel, J. Catal. 145 (1994) 335.
[18] K. Jothimurugesan, S.K. Gangwal, in: Proceedings of the US DOE Advanced Coal-Fired Power Systems’96 Review Meeting, Morgantown, WV, July 1996. [19] W. Wang, N. Padban, Z. Ye, A. Andersson, I. Bjerle, Ind. Eng. Chem. Res. 38 (1999) 4175. [20] M. Boudart, G. Djèga-Mariadassou, Kinetics of Heterogeneous Catalytic Reactions, Princeton University Press, Princeton, NJ, 1984, p. 98. [21] M.C.J. Bradford, P.E. Fanning, M.A. Vannice, J. Catal. 172 (1997) 479. [22] G. Ertl, M. Huber, J. Catal. 61 (1980) 537. [23] C.N. Satterfield, Heterogeneous Catalysis in Industrial Practice, McGraw-Hill, New York, 1991, p. 464. [24] P.B. Weisz, C.D. Prater, Adv. Catal. 6 (1954) 143. [25] W. Tsai, J.J. Vajo, W.H. Weinberg, J. Phys. Chem. 89 (1985) 4926. [26] R.W. McCabe, J. Catal. 79 (1983) 445. [27] J.-G. Choi, J. Catal. 182 (1999) 104.