Converting Wind Energy to Ammonia at Lower Pressure

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Converting Wind Energy to Ammonia at Lower Pressure Edward L. Cussler, Mahdi Malmali, Michael Reese, and Alon V. McCormick ACS Sustainable Chem. Eng., Just Accepted Manuscript • DOI: 10.1021/ acssuschemeng.7b03159 • Publication Date (Web): 07 Nov 2017 Downloaded from http://pubs.acs.org on November 13, 2017

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ACS Sustainable Chemistry & Engineering

Converting Wind Energy to Ammonia at Lower Pressure Mahdi Malmali1, Michael Reese2, Alon V. McCormick1, and E.L. Cussler1 1 Department of Chemical Engineering and Materials Science, 421 Washington Ave SE, University of Minnesota, Minneapolis MN 55455 USA 2 West Central Research and Outreach Center, 46352 State Highway 329, Morris, MN 56267 Corresponding author: [email protected] Abstract Renewable wind energy can be used to make ammonia. However, wind-generated ammonia costs about twice of that made from a traditional fossil-fuel driven process. To reduce the production cost, we replace the conventional ammonia condensation with a selective absorber containing metal halides, e.g. calcium chloride, operating at near synthesis temperatures. With this reaction-absorption process, ammonia can be synthesized at 20 bar from air, water, and wind-generated electricity, with rates comparable to the conventional process running at 150-300 bar. In our reaction-absorption process, the rate of ammonia synthesis is now controlled not by the chemical reaction, but largely by the pump used to recycle the unreacted gases. The results suggest an alternative route to distributed ammonia manufacture which can locally supply nitrogen fertilizer and also a method to capture stranded wind energy as a carbon-neutral liquid fuel.

Keywords: Reaction, Absorption, Energy Storage, Mechanism, Metal Halides, Mechanism

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Introduction Ammonia is one of the most important industrial chemicals in the world, the basis of chemical fertilizers and the “Green Revolution”.1 Without this fertilizer, three billion people in the world would be expected to starve.2 Ammonia is largely made via the Haber-Bosch process, a centuryold effort carefully and completely optimized at the large scale.3–5 This conventional process begins by catalytically reforming a fossil fuel in the presence of steam and air to produce a mixture of nitrogen, hydrogen, carbon monoxide, and carbon dioxide. This mixture is further reacted with steam through the well-known water-gas shift reaction to convert carbon monoxide to carbon dioxide. After the carbon dioxide is carefully removed with chemical scrubbing, the remaining mixture of nitrogen and hydrogen is reacted to produce the desired ammonia.3

While the reaction of hydrogen and nitrogen to form ammonia has a negative standard Gibbs free energy of change and hence could occur spontaneously, the rate is extremely slow. Overcoming this has led to a century of research, which continues today.6,7 The current solution even with the best available catalyst is to use high temperature to break the nitrogen-nitrogen triple bond and to reduce the amine adsorption energy to catalytic sites, and hence increase the rate of reaction. Such higher temperatures reduce the equilibrium reaction conversion, so that the existing process also requires high pressure (>150 bar) to increase the single pass conversion. This is the current state of this process, which is dangerous and hence very carefully operated in large, centralized facilities.3,8 Electricity generated from wind is a technology undergoing commercialization now.9 However, harvesting wind energy faces serious shortcomings. Wind energy is periodic, generated only when the wind blows. It is hard to store and to transport. Regions with high wind resources are typically in rural areas with limited populations and therefore, small electric loads. Wind farms located in these regions are at times, curtailed (shut down) due to constrained electric grids. The convergence of small local electric demand, extremely limited transmission capacities, and a high wind resource results in “stranded energy”.10 Battery storage is feasible but handicapped by low energy density and short life cycle, and so is impractical on the scale required.11


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Storage and transmission of wind-generated electricity is simplified if the electricity is converted into a carbon-neutral liquid fuel. Liquids are a core of current energy infrastructure based on fossil fuels. Liquid ammonia is an alternative to these fuels, with relatively high volumetric energy density. Ammonia serves as a carrier for hydrogen, which can be used in fuel cells.12 Ammonia can be used as a fuel in internal combustion engines, although this causes pollution unless optimized.13 Ammonia can also be used near where it is generated as a fertilizer. While this fertilizer is a key to the "Green Revolution",14 it is separate from using ammonia as a means of storing and transporting wind energy.

Wind-generated ammonia may be best produced in many small plants rather than in a few much larger plants.15 If the energy is made in multiple small plants, it can both utilize dispersed wind resources, and stabilize the aggregate electricity output. Such multiple small plants may benefit from existing ammonia storage infrastructure and handling experience in more remote agricultural areas. Ammonia may be collected and then shipped to population centers, like those in the northeast US, using existing pipelines.

Making small-scale ammonia from wind-generated electricity is feasible using the existing Haber-Bosch process.7 We are currently operating one such process in Morris, MN USA.16 This process does not use fossil fuels at all, but still requires feeds of pure nitrogen and pure hydrogen. It obtains the nitrogen from the pressure swing absorption of ambient air. It obtains the hydrogen from the electrolysis of water. The process, which uses the existing conventional catalyst, is 30,000 times smaller than a conventional ammonia plant, yet produces ammonia at about twice the cost of a conventional fossil fuel-based plant. This is expected: smaller plants routinely are more expensive per mass produced than large ones.

We can try to reduce the cost of small-scale wind generated ammonia in two ways. First, we can seek a new catalyst that operates at a lower temperature.17 Even though this is a good idea, a century of research shows that discovering such a catalyst is not trivial; thus it is not the focus here. A second way to reduce the process cost is to reduce the process pressure by selectively absorbing the ammonia almost immediately after it is synthesized. This is viable, as this paper shows. 3 ACS Paragon Plus Environment


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The small-scale conventional process is compared with our low-pressure approach in Figure 1. The conventional process, shown on the left, can be idealized as a reactor, a condenser, and a pump. The nitrogen and hydrogen flow through the catalytic reactor to make ammonia. Conversion is inhibited by the reverse reaction. The resulting gas mixture flows to a condenser, where liquid ammonia is partly removed. The unreacted gases and the ammonia residue are next recycled through the compressor to the reactor, which also is fed with more nitrogen and hydrogen to replace that consumed by the reaction. Our new process, shown schematically on the right of Figure 1, can also be idealized as three pieces of equipment: a reactor, an absorber, and a pump. It operates at the same high temperature as the conventional process, but at about 10 times lower pressure, and hence requires a smaller capital and operating investment. Ammonia partial pressure at lower reaction pressure is smaller, so the low-pressure process requires a very efficient separation. Absorption can provide such separation at high temperature. As before, gases are fed to the reactor; now, ammonia is more completely removed by absorption into solid ammine absorbents; and unreacted gases are recycled.

Figure 1 Conventional vs. Reaction‐Absorption Process. The condenser in the conventional process is replaced by an absorber, allowing operation at lower pressure, and reducing the need for heat integration.


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In this paper, we report characteristics of the low-pressure process and provide a roadmap for its further improvement. In our earlier work, we showed that the productivity of the entire reaction‐ absorption process is influenced not by the chemical kinetics, but by absorption and by recycle. The reaction-absorption process is faster because of the faster separation and recycle. To explore this in more detail, we use a simplified model for the rate of ammonia synthesis a, given in moles per time 16,18:

𝑎𝑎 = 𝑚𝑚𝑐𝑐𝑐𝑐𝑐𝑐 𝑟𝑟 =

𝑥𝑥 ∗ −𝑥𝑥0

𝜏𝜏𝑟𝑟𝑟𝑟𝑟𝑟 +𝜏𝜏𝑠𝑠𝑠𝑠𝑠𝑠 +𝜏𝜏𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟

(1)

where x* is the equilibrium mole fraction of ammonia at the reaction pressure and temperature, x 0 is the ammonia mole fraction at absorber outlet and actually in the gas in the absorber; and

τ rxn .n rxn , τsep .n sep , and τrecycle .n recycle (n i is moles of gas present in the unit “i”) are characteristic times of the chemical reaction, the separation, and the pump for recycling gases, respectively; r is the production rate given in moles per time per mass of catalyst. This equation includes a linearization of the chemical kinetics, so the characteristic times may vary both with pressure and temperature. Times which are large limit the overall synthesis rate; times which are relatively small have little effect. This equation successfully correlates data for our conventional smallscale process.16 The characteristic time of the reaction τ rxn .n rxn in this conventional small-scale process is largest, and the concentrations x 0 and x* are similar in size. Hence the rate a of the small-scale conventional process seems most strongly influenced by the chemical kinetics, represented by τ rxn , and by the reverse reaction, included as x*. It is less strongly affected by

τ sep and τrecycle . In the work reported below, we show that this is not the case for the new, lower pressure, reaction-absorption process. We also show how this new process can be further improved.

Experimental Materials Unless otherwise mentioned, all chemicals were ACS reagent and were used as purchased without further purification. Anhydrous, magnesium chloride (powder form, >%98 purity), calcium chloride (granular form, >93% purity), silica gel (technical grade), and 5 ACS Paragon Plus Environment


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aluminum oxide (>98% purity) were purchased from Sigma Aldrich (St. Louis, MO); Magnesium bromide hexahydrate (>99.5% purity) was purchased from Chem-Impex International Inc.

Synthesis of Supported Absorbents Supported metal halide sorbent materials were prepared using a liquid impregnation technique. Prior to impregnation, both 10 g of salt and 15 g of support were dried in a vacuum tube furnace at 450 °C with nitrogen sweep flow. Then both salt and support were added to a two-neck flask containing both 100 mL of ethanol (200 proof) and methanol. The contents of the flask were stirred, heated and fully refluxed for about 3 h with nitrogen purging. After the solvent was evaporated under vacuum, the residue was dried in a vacuum oven at 120 °C for 24 h.19

Reaction-Absorption Tests The schematic diagram of the reaction-absorption process is illustrated in Figure 2. A pre-reduced non-stoichiometric ferrous oxide (wustite) catalyst (AmoMax10, Clariant, Charlotte, NC) was used in these experiments. Catalyst particles ground to less than 1 mm particles were loaded into the reactor. The catalyst was activated under hydrogen for 48 hours from room temperature to 450 °C. The absorbent materials were also further dried after loading at 400 °C under nitrogen sweep flow. Both pure and supported salts were packed into a fixed bed column, and breakthrough curves were recorded using gas chromatography.19,20

We studied ammonia reaction-absorption in two different modes: batch reaction and fed-batch. In these tests, the apparatus was charged initially with hydrogen and nitrogen with a ratio of 3:1 until the desired pressure was achieved. In the batch experiments, the inlet valve was initially closed and the pump started to recirculate the gas mixture. In fed-batch tests (either reaction or reaction-separation), the inlet needle valve was left open. After starting these tests, more stoichiometric feed was injected by mass flow controllers (MFCs) once the system’s pressure dropped by less than 3%, thus sustaining the total pressure at the initial settings (e.g. 300 psi.) In batch experiments, system’s pressure was used as a means of evaluating the apparent production rate; whereas in fed-batch tests, the MFCs’ injections were used to calculate the ammonia synthesis rate. PID controllers were employed to control reactor, preheater, and 6 ACS Paragon Plus Environment

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absorber temperature. Reactor and preheater temperatures were set at 380 - 400 °C, and the absorber temperature was varied between 180-320 °C. LabView software (Austin, TX) was employed to control MFCs and to record pressure data. After each reaction-absorption test, the absorber was isolated and regenerated at 400 °C and atmospheric pressure with 20 cubic centimeters per minute of nitrogen sweep. Regenerations of one hour were found to be sufficient for regeneration.18

Figure 2. Reaction-Absorption Experimental Apparatus. This equipment was operated in either a batch or a fed-batch mode.

Results The results of this work are organized under three headings. First, we discuss the absorbents, focusing on those used by us and others in earlier work. These absorbents frequently have low accessible capacity and are unstable when used in the reaction-absorption process. Here we simply note these characteristics and defer the broad and complex topic of absorbent development to future papers. Second, we consider the variation of the ammonia synthesis rate with temperature and pressure. This rate varies linearly with pressure but less with temperature than expected from the chemical kinetics alone, suggesting that the rate limiting step is less due to chemistry than physics. Third, we show that the production rate varies sharply with the



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recycle rate of the unreacted gases. This observation provides a strategy for further increasing the rate of ammonia synthesis.

Absorbent Characteristics Metal halides have extraordinary potential capacity for ammonia storage.21,22 However, bulk metal halides are not stable during hot, multiple absorption‐ desorption cycles, and achieving the theoretical capacity for ammonia is difficult.19,23 Figure 3a shows breakthrough curves for a packed bed 10 cm long of 100 µm particles of pure magnesium chloride and magnesium bromide hydroxide. In case of magnesium chloride, the breakthroughs get shorter and shorter as the bed is repeatedly used and regenerated; in case of magnesium bromide hydroxide, the breakthrough curves are dispersed. These shorter breakthrough times mean that the capacities of the beds under these cyclic operating conditions are dropping. These capacity drops are apparently caused by the fusion of the pure absorbent, reducing the porosity and the surface area.

Metal halides in pure bulk form undergo crystal rearrangement during ammonia absorption and desorption cycles, which leads to gradual destruction of sorbent materials.12,24 In our earlier work we showed that granular form of calcium chloride is more suitable and stable for ammonia absorption-desorption;18 calcium chloride changes in the first few cycles. Then it remains unchanged, showing reproducible loading and unloading cycles. We have evaluated the performance of calcium chlorides for more than 100 cycles, as shown in Figure 3a, which led to reproducible capacities. However, this achieved only small ammonia loadings. Hence unsupported absorbent particles are not promising.19

Supporting the absorbents on inert porous materials like silica and alumina can both stabilize their performance and increase their apparent capacity. The data in Figure 3 illustrate these findings. In contrast, particles of supported magnesium chloride or magnesium bromide hydroxide, shown in Figure 3b, not only have a larger apparent capacity, but also exhibit reproducible breakthrough curves. Thus, the breakthrough time is both constant and larger for the same amounts of supported absorbents than for the unsupported ones.


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Figure 3. Ammonia Absorption Breakthrough Curves for Magnesium Chloride and Magnesium Bromide Hydroxide in Pure (a) and Supported (b) Forms. Bulk chloride and bromide do not show sharp repeatable breakthrough curves. Supported chlorides and bromides do display reproducible breakthroughs with higher effective capacity. (Absorption Temperature: 150 °C; Total Pressure: 1bar g; N 2 Flow: 50 SCCM; NH 3 Flow: 10 SCCM; Bulk MgCl 2 : 5 g; Bulk MgBr 2 : 3 g; Bulk CaCl 2 : 70 g; MgCl 2 –Al 2 O 3 and MgBr 2 (OH) n –Si: 10 g ) 9 ACS Paragon Plus Environment


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Synthesis Rate vs. Pressure and Temperature The pressures of batch reactions, a measure of the initial rate of ammonia production, are shown vs. time in Figure 4. The pressures, divided by their initial values to make the data more compact, show that the pressure change and hence the initial rate is roughly proportional to the initial pressure. The slope of each curve provides a measure of the rate of ammonia synthesis at any time. At small times, the production rates remain constant, confirming that absorber is capable of removing ammonia produced in the reactor and thus maintaining a steady production rate. At longer times (~5000 s), the drop of the rates may be either due to the drop in the operating pressure or to partial saturation of the absorber.

Figure 4. Chemical Reaction Rate vs. Pressure. In these recirculating batch experiments, the production rate is roughly proportional to the pressure change. These data are for CaCl 2 at 180 °C and 1.55 mmol/s of recycle.


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The synthesis rate of fed-batch experiments is also dependent on pressure, as shown by the three types of experiments shown in Figure 5. The first type, at 20 bar and with reaction temperature of 400 °C, has neither a condenser nor an absorber. As expected, the rates of this process, shown as crosses in the figure, are very low because without separation, the fast forward reaction is quickly inhibited by the reverse reaction. This process is ineffective both at high pressure and at low pressure, so has no commercial value.

The second type of constant pressure process, shown as the dotted line, presents results from our small conventional pilot plant operating with a condenser. The reactor in this steady-state process runs at 120 bar and 440 °C, using the same catalyst with the same activity as that used commercially.16 The condenser, running at -20 °C, yields liquid ammonia and a gas mixture nearly saturated with ammonia. The condenser’s performance is consistent with that predicted from correlations given in the literature. This process is effective at making ammonia, but at a cost about twice that of the much larger commercial alternative.

Figure 5. Ammonia Production at Constant Pressure. In the fed-batch experiments near 20 bar, the production rates with absorbent (points) are higher than in the conventional plant with 11 ACS Paragon Plus Environment


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condensation (dotted line), even though the conventional plant is running at 120 bar. The rate without absorbent or condensation (crosses) is much smaller because of reverse reaction. The third type of process, shown as the data points, is the laboratory-scale reaction-absorption system, run in this case with a calcium chloride absorbent. This fed-batch process operates at a reaction temperature at 400 °C, close to the large conventional and the small plants, but at a lower pressure of 20 bar. The absorber, packed with 3-7 mm calcium chloride particles, operates at temperatures of 180 or 320 °C. Note that results reported for the reaction-absorption process are being run only at times which do not require absorbent regeneration. Even though the lab process is running at six times lower pressure, it gives production rates close to those of the small-scale conventional process. Because it is at least as fast and operates at a lower pressure, the lab process should result in lower capital and operating expenses associated with containing and compressing the gases, exchanging heat, and chilling a compressor. Both the results in Figure 4 and in Figure 5 show an ammonia synthesis rate which is pressure-dependent. This is consistent with Equation 1.

The ammonia production rate also varies with temperature, as shown in Figure 6 for more batch experiments. As in Figure 4, the pressure divided by the original pressure is plotted vs. time to make the data more compact. The rate of pressure change does not vary much with temperature; also, the final conversion is greater for the lower temperature. This is consistent with a rate controlling step for these experiments which is not the chemical kinetics. To reinforce this point, the figure also plots as the dashed line the pressure change expected from literature kinetics measurements of the forward chemical reaction, with no inhibition by the reverse reaction. It plots as the solid line the pressure expected from literature measurements of the chemical kinetics of both the forward and reverse reactions. Both these lines show that the separation is dramatically affecting the results, and that the overall rate of the reaction-absorption process is not dramatically influenced by temperature and hence by the chemical kinetics. In other words, Figure 6 implies that τ rxn .n rxn is smaller than τ sep .n sep and τ recycle .n recycle . In passing, we note that this process is still showing significant conversion even at pressures less than one bar, a point which we will discuss more later.


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Figure 6. Pressure Changes vs. Temperature. The pressure changes with time at an initial pressure of 30 bar compliment the results at various pressures and 180 oC shown in Figure 4. The dashed line represents the chemical kinetics without any reverse reaction, and so illustrates the kinetics possible with extremely fast absorption and recycle, while the dashed-dot lines represents a reaction without a separation.

Synthesis Rate vs. Recycle Flow So far, we have discussed the nature of the absorbent used and the ammonia synthesis rate as a function of pressure and temperature. Finally, we consider the variation of synthesis rate with the recycle flow. Equation 1 predicts that the reciprocal of synthesis rate (1/a) should be proportional to the time for recycle τ recycle .n recycle . Figure 7 gives an alternative form of this plot, as the reciprocal of the rate per catalyst mass (1/r=m cat /a) vs. the reciprocal of the recycle flow (1/m recycle ). The expected variation is observed. This is true both



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for the experiments with unsupported calcium chloride and with magnesium chloride stabilized by alumina. It is true for experiments both after 500 s., and after 1000 s.

The results in Figure 7 suggest that the rate of the reaction-absorption process shown on the right-hand side of Figure 1 is controlled largely not by the chemical kinetics or by the rates of absorption, but by the rates of recycle of unreacted nitrogen and hydrogen. In other words, these results imply that the times for reaction τ rxn .n rxn and absorption τ sep .n sep are smaller than the time for recycle τ recycle .n recycle . The basis of this conclusion is discussed in the final section of this paper.

Figure 7. Production Rate vs. Recycle Flow. The reciprocal of the initial production rate is proportional to the reciprocal of the recycle flow. The slope of these data reflects resistance of recycle, and the intercept includes constraints of the kinetics of chemistry and of absorption. Pressure: 55 bar; absorber temperature: 180 °C.


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Discussion This paper explores an alternative process for making ammonia. The process uses wind energy to make electricity, which in turn is used to make nitrogen from the separation of air, and hydrogen from the electrolysis of water. These gases are combined over a Haber process catalyst to make ammonia. However, when the wind-based process is a small-scale clone of that used in conventional, fossil-fuel based ammonia manufacture, the ammonia made from wind costs more than that currently made from fossil fuels. If wind-based ammonia is to become commercially attractive, it must be cheaper. Wind is relatively cheaper than any other type of fuel. Hence, the cost of making ammonia cannot be reduced by replacing the energy source. The cost can be reduced if the equipment for carrying out this synthesis requires less capital, and also the operating such process (compression, exchanging heat, and separation) is cheaper.25 Such a simple, low capital process is our goal.

To reach this goal, we have begun studying a low-pressure ammonia synthesis which separates ammonia not with the current condenser but with an ammonia-selective absorption bed. Such a low-pressure process can potentially cost less because the upfront capital investment and day-to-day operating costs are lower than that of the conventional process. This paper explores the mechanism of this process by identifying its rate-controlling step. Once this step is known, we will be able to optimize the process design. The data given above begin this search. Those in Figure 3 show that the ammonia synthesis rate varies with pressure, as expected from the linearized analysis in Equation 1. The data in Figure 4 illustrate how the production rate can still be maintained by removing ammonia as soon after it is formed, that is, by keeping (x*-x o ) in Equation 1 large even if the values of x* and x o are smaller. The results in Figure 5 show that the overall synthesis rate still rises at higher pressure, so that any reductions in pressure are being overcome with increases in overall kinetics. The data in Figure 6 show that the reaction rate varies less with temperature than expected from the chemical kinetics alone, so chemical kinetics



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is not the rate controlling step: τ rxn .n rxn in Equation 1 must be smaller than other characteristic times.

But the important results are illustrated in Figure 7, which imply that the rate limiting step in the low-pressure process is not the rate of chemical kinetics, nor the rate of ammonia separation, but just the rate at which the gases are recycled. This graph plots the reciprocal of the production rate vs. the reciprocal of the recycle flow. If Equation 1 is a good model of the low-pressure process, this graph should be linear, and it is. More specifically, the intercept on this graph is the resistance to synthesis of both the chemistry and the separation; the slope gives the influence of the recycle. At slow flows, the effect of the recycle is around 80% of the total resistance. At high flow, the recycle rate has a smaller affect, but under the experiments made so far, it is always about half the total resistance, implying that τ recycle .n recycle is about equal to τ rxn .n rxn plus

τ sep .n sep . To increase the overall ammonia synthesis rate for the current reaction-absorption process, we need not to seek a new catalyst nor a better absorber; we need only to buy a bigger pump to recycle the gases even faster.

If true, this conclusion gives an easy way to increase ammonia synthesis rate, and hence to reduce the capital cost per mass of ammonia made. This is true even at the lower pressures used. But is this conclusion valid? To explore this further, we estimate each of the times in Equation 1, and show that they are similar to those observed experimentally.

The Rate Limiting Step We begin by estimating the time constant for the chemical reaction τ rxn for six of the kinetic rate equations reported in the literature. To do so, we consider the models for ammonia synthesis kinetics suggested by Temkin and Pyzhev 26, Nielsen et al. 27, Stoltz and Norskov 28–30, and Bowker et al. 31–33. The model reported by Temkin and Pyzhev (case 1) asserts:

𝑟𝑟 = 𝑘𝑘1𝑜𝑜

3 𝑃𝑃𝑁𝑁2 𝑃𝑃𝐻𝐻 2 2 𝑃𝑃𝑁𝑁𝑁𝑁 3

− 𝑘𝑘2𝑜𝑜

2 𝑃𝑃𝑁𝑁𝑁𝑁 3 3 𝑃𝑃𝐻𝐻 2

The model used by Nielsen et al. (cases 2-4) is: 16 ACS Paragon Plus Environment

(2)

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𝑟𝑟 =

𝑘𝑘1𝑜𝑜 𝑃𝑃𝑁𝑁2 −𝑘𝑘2𝑜𝑜

2 𝑃𝑃𝑁𝑁𝑁𝑁 3� 3 𝑃𝑃𝐻𝐻 3

𝑃𝑃𝑁𝑁𝑁𝑁3 �1+𝐾𝐾3 �𝑃𝑃 𝜔𝜔 � 𝐻𝐻2

(3)

2𝛼𝛼

where the adjusted rate constants are given in Table 1.34

Table 1. Rate constants for cases 2-4.

Case # 2 3 4

𝐾𝐾30

𝐸𝐸3

𝑘𝑘20

𝐸𝐸2

𝛼𝛼

𝜔𝜔

3.07 × 10−2

−19361

2.12 × 1013

17249 0.640 1.564

2.96 × 10−6

−29241

4.60 × 108

4854

2.94 × 10−4

−24052

7.89 × 1010

𝐴𝐴 = 𝐸𝐸2 − 2𝛼𝛼𝐸𝐸3 42300

11171 0.654 1.523

42600

0.661 1.454

43500

The model developed by Sehested et al. 35 (cases 5-6) reports the rate equation:

𝑟𝑟 = 2𝑁𝑁𝑆𝑆 𝐾𝐾1 𝑘𝑘2 𝜃𝜃∗2 �𝑃𝑃𝑁𝑁2 −

2 𝑃𝑃𝑁𝑁𝑁𝑁 𝑃𝑃0.5 𝑃𝑃 3 � �1 + 𝑁𝑁𝑁𝑁3� 1.5 + 𝐻𝐻2 �𝐾𝐾 � � 3 𝑃𝑃𝐻𝐻3 𝐾𝐾𝑒𝑒𝑒𝑒 𝑃𝑃𝐻𝐻2 𝐾𝐾𝑎𝑎 𝑏𝑏

where the values of the rate constants suggested by different investigators are listed in Table 2. These values presume the number of active sites on the catalyst is 52 the value for the catalyst used in our experiments.

𝜇𝜇𝜇𝜇𝜇𝜇𝜇𝜇 35 , 𝑔𝑔

which is similar to

Table 2. Rate constants proposed for cases 5 and 6. Case # 5

6

𝐾𝐾𝑎𝑎

5.15 × 10−3 exp�−31.7�𝑅𝑅𝑅𝑅�

5.54 × 10−2 exp�−53.0�𝑅𝑅𝑅𝑅�

2 𝐾𝐾1 𝑘𝑘2

116 exp�14.6�𝑅𝑅𝑅𝑅� 974 exp�14.7�𝑅𝑅𝑅𝑅�

𝐸𝐸𝑏𝑏

2.64 × 109 exp�−155�𝑅𝑅𝑅𝑅�

5.26 × 1012 exp�−199�𝑅𝑅𝑅𝑅�

−𝐸𝐸𝑏𝑏� −1 𝑅𝑅𝑅𝑅� and 𝐸𝐸𝑏𝑏 = 48 𝑘𝑘𝑘𝑘 𝑚𝑚𝑚𝑚𝑚𝑚

𝐾𝐾𝑏𝑏 = 2.16 × 103 exp �


Reference Stoltz and Norskov 28–30 Bowker et al. 31– 33


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−𝐸𝐸 𝐾𝐾𝑒𝑒𝑒𝑒 = 2.03 × 10−12 exp � 𝑒𝑒𝑒𝑒�𝑅𝑅𝑅𝑅� and 𝐸𝐸𝑒𝑒𝑒𝑒 = −101.6 𝑘𝑘𝑘𝑘 𝑚𝑚𝑚𝑚𝑚𝑚 −1 Using these equations and the values of the rate constants given in these tables, we can now calculate the contribution to the intercept in Figure 7 of the catalyst mass m cat times the reaction time τ rxn for the conditions used in our experiments. The details of the calculations are summarized for the six cases in Table 3. Table 3. Calculated intercepts due to reaction in Figure 7. The estimates are from by different models reviewed in the text.

Case #

1/r [g cat s mol-1]

1

0

2

400

3

1200

4

1800

5

270

6

32

The models predict a contribution to the intercept of less than 2000, a fraction of the observed intercept of 10000. (The model of Temkin and Pyzhev is extreme: it predicts a contribution to the intercept of zero because Equation 2 predicts an infinitely fast rate when the ammonia concentration is zero, which would be approached if the absorption and the recycle were both infinitely fast.) These six predictions imply that the chemical kinetics has only a small effect of the rate of ammonia synthesis in the reaction-absorption process, which was our conclusion from the very different experiments reported above.

The remainder of the intercept, with a value around 8000 g cat s/mol, is apparently due to the separation by absorption. This remainder equals m cat /k c A, that is, the mass of catalyst m cat divided by the product of the mass transfer coefficient k c and the total absorbent area A c . Regrettably, we have at present no separate measurements of k c and A c ; we do not know either 18 ACS Paragon Plus Environment

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the mechanism responsible for mass transfer or the interfacial area which is most relevant. We can see that the time for absorption τ sep can be decreased by using more absorbent or by adapting the same volume of an absorbent which has a greater surface area per volume than the one currently being used. This supplies a simple strategy for reducing this part of the intercept in Figure 7, and hence increasing the overall rate of ammonia synthesis.

In addition to the intercept, we can estimate the slope of the results in Figure 7, which should equal the mass of catalyst divided by the molar flow rate m. The experimental value of this slope is 5 g cat s/mol, about three times smaller than our direct experimental value of around 15 g cat s/mol. This discrepancy may partly be due to our difficulty of accurately measuring a flow at high temperature and pressure, which deviates from our calibration measurements at ambient conditions. This calibration may also be responsible for the difference between the results of the two different absorbing salts. Nonetheless, the results still follow the prediction of Equation 1, and hence support the linearized model for the entire process.

Further Improvements These considerations suggest how the low-pressure process can be improved beyond what we have achieved to date. These improvements have two parts: the design of a better process based on what we now know; and a still better process using science we have not yet discovered. These are discussed sequentially.

The process based on what we now know involves improvements to the reactor, the absorber, and the pump. The reactor should continue to operate around 400 °C. While lower temperatures could provide smaller problems of energy management, they produce too great a drop in chemical kinetics. We can reduce the pressure from 150-300 bar to less than 20 bar. While this reduces the chemical kinetics, it also reduces the capital cost of the equipment and the size of the compressor, thus cutting the equipment cost. Operating cost, specially cost of compression will be much smaller.



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At present, we would base the absorber on magnesium bromide hydroxide supported on silica. Of the absorbents studied to date, this has the greatest stability and the largest effective capacity. We hope to operate this absorber as close to the reactor temperature of 400 °C as possible, and to regenerate with only reductions in pressure, and without increases in temperature. At present, we are not completely sure if this is possible, but we believe that it is within reach. We will increase the size of the pump by about a factor of five. These changes to reactor, absorber and pump should make the time constants about equal, and give yield an ammonia synthesis rate about four times greater than currently observed.

Some readers will note that we do not plan to try to improve the catalyst, a sensible but difficult goal. This continues to be an active topic of research by others. However, under the current conditions, the catalyst limits the ammonia synthesis rate by less than 10% of the total, so if we were to improve the catalyst by a factor of two, we would not really gain much. The advantage to improving the catalyst, of course, is that it could potentially increase the equilibrium conversion. It could also increase conversion at lower temperature. Still, we feel the past struggles in this direction make our odds of dramatic success relatively small.7 We now are investigating the practical value of our current results.

At the same time, we are intrigued by the possibility of still higher synthesis rates at low pressure by integrating the reactor and the separator. After all, if each active site on the catalyst were within, say, 100 µm of an active site for absorption, we can potentially get still higher rates at still lower pressures. We are charmed by this vision, but we understand its enormous problems. For example, how would the absorbent be regenerated? How would we minimize catalyst poisoning? Thus we will focus on examining the implications of what we report above, even while we dream of more advances.

Acknowledgments This work was primarily supported by the US Department of Energy (ARPA‐E, USDOE / DEAR0000804), by the Minnesota Environment and Natural Resources Trust Fund (LCCMR, / ML 2015, CH 76, SEC 2, SUBD 07A), and by the MnDRIVE initiative of the University of Minnesota 20 ACS Paragon Plus Environment

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(MNT11.) Other support came from by the Institute for Renewable Energy and the Environment (IREE, RO-0001-12) at the University of Minnesota, and by the Dreyfus Foundation. Michael Ho helped with some calculations.

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10.1006/jcat.1999.2628.


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For Table of Contents Use Only Renewable ammonia from stranded wind can be used to generate energy. It can produced via low pressure reaction-absorption process, which is not controlled by reaction but by rate of recycling.


Converting Wind Energy to Ammonia at Lower Pressure

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