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Ammonia Synthesis at Reduced Pressure via Reactive Separation Mahdi Malmali, Yong-Ming Wei, Alon V. McCormick, and Edward L. Cussler Ind. Eng. Chem. Res., Just Accepted Manuscript • DOI: 10.1021/acs.iecr.6b01880 • Publication Date (Web): 25 Jul 2016 Downloaded from http://pubs.acs.org on August 1, 2016

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Industrial & Engineering Chemistry Research

Ammonia Synthesis at Reduced Pressure via Reactive Separation Mahdi Malmali†, Yongming Wei‡, Alon McCormick†, Edward L. Cussler†* †

Department of Chemical Engineering and Materials Science, University of Minnesota, 421 Washington Ave SE #151, Minneapolis, Minnesota 55455, United States

‡

Chemical Engineering Research Center, East China University of Science and Technology, 130 Meilong Rd., Shanghai 200237, China

Abstract Ammonia is normally made at high temperature and pressure using a promoted iron catalyst. The high temperature is needed to get fast kinetics; the high pressure is used to ensure high conversion. Alternatively, ammonia can be made at high temperature but lower pressure if the product ammonia is rapidly separated. Here we have systematically studied the effect of temperature and pressure on the rates of reaction. Then we have qualitatively investigated the absorptive separation of ammonia using calcium chloride in a reaction-separation process. Rapid separation reduces the constraint of reversible reaction and enables us to obtain appropriate reaction rates at relatively lower pressure. The effect of different operating conditions - reaction temperature, pressure, absorption temperature and gas transport - on production rates is carefully measured, and this elucidates the potential and the limits of this type of low pressure ammonia synthesis.

Keywords: Reaction-Separation, Ammonia, Absorption, Reaction Rate Constants, Metal Amine Complex

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Introduction

Those in the global chemical enterprise now agree that the future of the chemical industry depends on becoming more sustainable.1 Such a goal includes developing energy sources that are not based on non-renewable fossil fuels, and that do not release large quantities of carbon dioxide. This goal implies improving the efficiency of energy collected from the sun and the wind. Solar energy is continuing to get cheaper, but is still more expensive than fossil fuels. Wind energy is more immediately attractive, and is being rapidly developed in areas of high population density, especially in northern Europe.2

However, both wind and solar energy are periodic, and hence may not be available at times of highest demand. No one wants light only in the middle of a sunny day; no one wants power only when the wind is blowing. Moreover, wind and solar productions are distributed geographically, with practical limitations in supplying the electrical grid in remote locations. Thus both of these resources must be coupled to methods of energy storage, especially as liquid fuels. Methanol and ammonia are two such fuels. Methanol is attractive because it meshes well with the existing infrastructure for distributing fuel. Ammonia is interesting because it can be made from totally renewable resources: nitrogen produced from pressure swing adsorption of air, hydrogen produced from electrolysis of water, and electricity obtained from stranded solar arrays or wind


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turbines. Thus, a major advantage of ammonia over methanol as a fuel is that ammonia does not require a carbon source for the synthesis.

We are especially attracted by ammonia because much of the potential wind power is far from areas of high population. As a result, wind is sometimes called “stranded energy.”3 At the same time, much of the stranded wind energy is where ammonia based fertilizer is needed. Thus by working on ammonia, we are investigating a possible chemical synthesis valuable both as a carbon-neutral liquid fuel4,5 and as a fertilizer.6

We have begun to investigate making ammonia from wind using a small ammonia plant operating at the West Central Research and Outreach Center (WCROC) in Morris MN.6 This plant lets us explore what parts of a conventional process work well, and what parts are potential problems. The plant uses a reactor, a condenser, and a compressor, all of which are scaled down parallel to a conventional large-scale plant.

The reaction rates observed in the plant are

consistent with parallel laboratory experiments with the same catalyst, and with values from the literature.7–10 Analysis of this process gives three characteristic times, one for the reaction, one for the condensation, and one for the pump recycling unreacted gases.

As the process is

currently operated, the reaction time is largest, so that the small process is currently controlled by the chemical kinetics. However, the analysis also shows how process changes could increase productivity until one of the other characteristic times became largest, and hence the limit of production. Thus this plant supplies a test for process changes that may make sense for a small-



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scale plant, even when they are known to have limited effectiveness in a large fossil fuel based synthesis.

In this paper, we explore the potential of process changes, to reduce the pressure needed for production. In seeking these process changes, we have decided not to focus on the reaction chemistry, feeling that this has been so carefully studied over the last century that the chances of big improvements at any scale are small.11–14 Instead, we are investigating three other aspects of this process. First, we are investigating synthesis at lower pressure, because such pressure could reduce process complexity and make a small-scale process cheaper and less hazardous. Second, we are investigating ammonia-selective absorbents to reduce the need for recycle. Third, we are beginning to explore processes which avoid recycle altogether.

In this paper, we limit the discussion just to the first aspect, studies of lower pressure synthesis. After a summary of the process itself, we report measurements of chemical kinetics to show that at low pressures and a range of temperatures, we get the same results as others working under these conditions. We then report the measurements of ammonia production in a reactionseparation process. Calcium chloride absorbents are employed, where absorption occurs not just on the surface of the solid, but also by diffusion into the solid. Absorption and desorption15 of ammonia on chlorides of alkaline-earth metals have been studied in numerous applications, such as chemical heat pumps,16 ammonia storage,4,17,18 indirect hydrogen storage,19,20 and enhanced ammonia synthesis.21 We use the measured combination of reaction and absorption to estimate the feasibility of ammonia synthesis at lower pressures.


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Overall Production Rate: Here, we employ a model that we developed in our earlier work6 for the reaction-separation process with recycle, which guides our analysis of our experiments. This theory helps us to understand the complex behavior of reaction-separation process, and we feel that this theory – specifically the time constants of the reaction, separation and transport – makes it easier to compare the findings of our reaction-separation process. In following, we are giving more details about the model that eases the comprehension of the theory and the model. More details on this model can be found elsewhere.6

In this simplified model, we assume that the process is in steady state, and that the reactor and absorber are each well mixed. (The steady state assumption is accurate only at the beginning of our experimental test.) The well mixed assumption implies that average nitrogen concentration in the reactor equals the nitrogen concentration at the exit.22 While this is untrue if most of the nitrogen entering the reactor reacts in a single pass, it is much more nearly true when only a small fraction of the nitrogen reacts per pass, which will be the case here. However, this will not interfere with our findings, as will be discussed later. The model predicts the rate of ammonia production a:

=

∗

∗

(1)

where ∗ and are the mole fractions of ammonia at reaction equilibrium and in the absorber,

respectively; and are the linearized chemical reaction rate constant and the absorption



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mass transfer coefficient, respectively; Pt is the operating pressure; S is the surface area of the absorber; and m is the total molar flow rate. In this result, the term 1⁄ ! is the resistance to

ammonia production due to the chemical reaction. Similarly, 1⁄ "# $! and 1 − ∗ !⁄& are the resistances of the absorber and of the recycle loop, respectively.

In this model the absorber’s performance is assumed to behave as a first-order rate process. In reality, before the breakthrough of the absorber’s packed bed, we have a perfect separation of ammonia from the gas mixture. Hence before breakthrough, the absorption resistance is effectively zero. After breakthrough, though, the production rate can become dominated by the raising absorption resistance.

The presence of three resistances in series is observed in many rate processes, where the amount produced is proportional to the overall driving force divided by the total resistance.23 The total resistance is the sum of the resistances of reaction, absorption, and recycle. Phrased in other terms, the total resistance is a harmonic average of the speeds of these three steps, so that the slowest speed has the biggest effect on the rate of ammonia production. If the reactor is operating at too low a temperature, the reactor resistance will be dominating; if the absorber is undersized, the absorption resistance can be largest; and if the recycle rate is too low to take full advantage of the other unit operations, the recycle resistance will be most important. We will use this model later in this paper to discuss the experiments, and hence to explore how the reactionseparation process can be made more productive.


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Linearized Chemical Reaction Rate Constant: Before we make this analysis, we want to review how the linearized reaction rate constant can be estimated from earlier studies. This reaction rate is most often correlated using the Temkin-Pyzhev8,9,24 equation:

' = (

.. )*+ ), +

)

− /

)

.. ), +

(2)

where k1 and k2 are forward and reverse reaction rate constants, respectively; and "0+ , "1+ , and

" are the partial pressures of nitrogen, hydrogen and ammonia, respectively. This equation is rewritten by defining the following variable:

2=

) ∗ ) )*+ ∗

(3)

Linearization using the Taylor series for small values of X simplifies this:

' = (

.. )*+ ∗ ), +

)

∗

∗

4

35 +

)* + ∗ )

∗

7 2 − 2 ∗ ! + /

) ∗

.. ∗ ), +

)* ∗

(

3 ) +∗ − 57 2 − 2 ∗ ! = 8 2 − 2 ∗ !

(4)

where 8 has the dimension of moles of ammonia per catalyst volume per time. The 8 is related to kR because: 2=

) ∗ ) )* + ∗

= 4 :∗ −

∗< ;(

+

=

(5)



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Hence, when the conversion is close to zero, which is the case in our analysis,

8 ≈ 4

(6)

where 8 is the corrected linearized reaction rate constant after change of variables. Equation 4

shows that the linearized reaction rate constants 8 and obtained here are functions of

concentration. When approaches ∗ , 8 falls to zero; but when xA approaches zero, 8 goes to its maximum value, representing only the forward reaction rate.

Experimental Materials: Anhydrous CaCl2 in granular form (>7 mm) with 93% purity (CAS#: 10043-52-4, Lot #: SLBL2770V) was purchased from Sigma Aldrich (St. Louis, MO). The reactant gases N2 and H2 with ultrahigh purity were purchased from Matheson (New Brighton, MN). We employed a pre-reduced non-stoichiometric ferrous oxide catalyst (wustite) with promoters, AmoMax-10 RS (Clariant, Charlotte, NC). The catalyst is provided in irregular shape granules, with the nominal size range of 1.5 – 3 mm, and is stabilized with an oxygen-rich protective layer.

Apparatus: The experimental apparatus, a schematic diagram of which is shown in Figure 1, was built using Swagelok (Chaska, MN) 316 stainless steel tubing and fittings of 6 mm inner


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diameter. The reactor was 0.15 m long. Catalyst particles were ground from their produced size until smaller (

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