Kinetics and Reaction Engineering Aspects of Syngas Production by the Heterogeneously Catalysed Reverse Water Gas Shift Reaction

1 Kinetics and Reaction Engineering Aspects of Syngas Production by the Heterogeneously Catalysed Reverse Water Gas Shift Reaction 1 Von der Fakultät für Angewandte Naturwissenschaften der Universität Bayreuth zur Erlangung der Würde eines Doktor-Ingenieurs (Dr.-Ing.) genehmigte Dissertation 1 vorgelegte von M.Tech. Unde Rajabhau Bajirao aus Undewadi (Indien) Erstgutachter: Prof. Dr.-Ing. Andreas Jess Zweitgutachter: Prof. Dr. rer. nat. Peter Wasserscheid Tag der mündlichen Prüfung 11. June 2012 Lehrstuhl für Chemische Verfahrenstechnik Universität Bayreuth 2012 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 Acknowledgements First of all I would like to express my sincerely thank to Prof. Dr. Andreas Jess for giving me the opportunity to work in this group, for making my wish become realistic, for his excellent guidance and support, for the enthusiasm in supervision and tiredless correction of my dissertation. His advice and insight into things throughout my doctoral project have been invaluable. My thanks go to Prof. Dr. Peter Wasserscheid for agreeing co-referee to my thesis, and also to Prof. Dr. Ruth Freitag and Prof. Dr. Ralf Moos for accepting to be in my examination committee. I thank to Dr. Christoph Kern for his support in the modelling, for providing valuable suggestions, and for the correction of my dissertation and also to Dr. Wolfgang Korth for correction of my dissertation and helpful comments as well as many useful discussions. I also thanks to Dr. Leonid Datsevich for his fruitful discussions, Mr. Jörg Gerchau for much help of experimental set-up and the operation of instruments, and for the help of solving computer problems, Mrs. Birgit Brunner for her help and support during my research work. I thank Secretary Mrs. Rita Pannek for her administrative assistances. I thank my colleagues from Chair of Chemical Engineering Johannes Thiessen and Lisa Schilder for a good time of sharing the office with them and for their help, Florian Heym, Amadeus Rose, Anne Piegsa, Stefan Fritz, Stephan Aschauer, Philipp Kaiser, Peter Fremerey and Susanne Fritschi for their help, support and providing pleasant and friendly research atmosphere over the years. I thank to all my friends and my country-mates in Bayreuth for their love, supports and encouragement. I am forever indebted to my parents, parents in law, family members and friends for their love, prayers, supports and encouragement. Finally, I thank to my lovely wife, Sudha, for her support, understanding, encouragement and endless patience. 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 Parts of this work were already published: 1 1. A. Jess, P. Kaiser, C. Kern, R. B. Unde and C. von Olshausen, Considerations concerning the Energy Demand and Energy Mix for Global Welfare and Stable Ecosystems. Chemie Ingenieur Technik 83, 1777–1791 (2011). 2. R. B. Unde, C. Kern and A. Jess, High temperature CO 2 hydrogenation over Ni Catalyst. 8 th European Congress of Chemical Engineering, ProcessNet, Berlin, September 25–29 (2011). 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 i Table of Contents List of symbols ....................................................................................................................... v 1 1. 1 Introduction ................................................................................................................. 1 1 2. 1 Basic theory and background of the work ................................................................ 3 1 2.1 Heterogeneous catalysis ............................................................................................... 3 1 Diffusion processes in heterogeneous catalysis ...................................................... 3 1 2.1.1 Influence of internal and external mass transport on heterogeneously catalysed 2.1.2 reactions ................................................................................................................. 5 1 Catalyst deactivation phenomena .......................................................................... 13 1 2.1.3 2.2 Utilization of CO 2 for production of chemicals and fuels .......................................... 16 1 2.3 Concept of production of liquid fuels from CO 2 via reverse water gas shift (RWGS) and Fischer Tropsch synthesis (FTS) ........................................................................ 20 1 2.4 Thermodynamics of CO 2 conversion (reverse water gas shift and methanation) ...... 27 1 2.5 Kinetics of reverse water gas shift reaction and methanation .................................... 32 1 3. 1 Objective and scope of the work .............................................................................. 39 1 4. 1 Experimental method and data analysis ................................................................. 41 1 4.1 Experimental setup ..................................................................................................... 41 1 4.2 Experimental procedure ............................................................................................. 43 1 4.3 Catalyst characterisation ............................................................................................ 45 1 4.4 Evaluation of the experimental data ........................................................................... 47 1 CO 2 hydrogenation reaction .................................................................................. 47 1 4.4.1 4.4.1.1 Conversion of CO 2 and yield of CO and CH 4 .............................................. 47 1 4.4.1.2 Determination of the intrinsic kinetic parameters ........................................ 48 1 4.4.1.3 Internal mass transport calculations ............................................................. 49 1 4.4.1.4 Calculation of external mass transport limitations ...................................... 51 1 CO hydrogenation (methanation) reaction ............................................................ 52 1 4.4.2 4.4.2.1 Conversion of CO and yield of CO 2 and CH 4 .............................................. 52 1 4.4.2.2 Kinetic analysis and mass transport calculations ......................................... 52 1 Water gas shift (WGS) reaction ............................................................................ 53 1 4.4.3 4.4.3.1 Conversion of CO and yield of CO 2 and CH 4 .............................................. 53 1 4.4.3.2 Kinetic analysis and mass transport calculations ......................................... 53 1 5. 1 Results and discussion ............................................................................................... 55 1 viii Dimensionless numbers Nu Nusselt number Pr Prandtl number 9A Reynolds number 3B Schmidt number 3 C Sherwood number Abbreviations CCC carbon capture and conversion FTS Fischer Tropsch synthesis HTE high temperature electrolysis of water MFC mass flow controller NTP normal temperature and pressure (20 °C and 1 atm) RWGS reverse water gas shift toe tonnes of oil equivalent, 1 toe = 42 GJ TPR temperature programmed reduction TPS technical photosynthesis WGS water gas shift XRD x-ray diffraction Introduction 1 1. Introduction Carbon dioxide is the most important greenhouse gas being emitted into the atmosphere from fossil fuel combustion and other anthropogenic activities. According to the National Oceanic and Atmospheric Administration (NOAA), the CO 2 concentration in the atmosphere was 392 ppm in February 2011 [1] which is very high compared to 379 ppm in 2005 and far from the natural range of the last 650,000 years [2]. The annual emissions of CO 2 have grown between 1970 and 2004 by about 80% from 21 to 39 gigatonnes [2]. The growing concentration of carbon dioxide in the atmosphere increased the impact on the environment such as global warming and forcing a climate change. The main greenhouse gases are water vapour (H 2 O), carbon dioxide (CO 2 ), methane (CH 4 ), nitrous oxide (N 2 O), hydrofluorocarbons (HFCs), perfluorocarbons (PFCs), and sulphur hexafluoride (SF 6 ), but CO 2 contributes about 77% to the world’s greenhouse gas emissions (excluding water vapour) into the atmosphere in 2004 [2]. Alternative energy sources are not only important with regard to global warming but also to maintain the rising cost of crude oil. Both aspects have motivated researchers to look for solutions to reduce the greenhouse gas emission and/or its utilization. The mitigation of greenhouse gas emissions is also an interesting challenge in exploring new concepts and new opportunities for catalysis and industrial chemistry. Recently various carbon capture and storage technologies (CCS) are developed for a significant reduction of CO 2 emissions into the atmosphere [3]. These technologies are mainly suitable to capture CO 2 from large industrial sources such as fossil fuel fired power plants. But only with these efforts it is not possible to reduce and control CO 2 emissions [4], and the availability of sufficient storage capacity to capture carbon dioxide is also still an open question. However, very little attention has been paid by industry and academia in utilization of CO 2 because of its high thermodynamic stability. For the reduction of greenhouse gas emission various strategies have been suggested that includes promoting energy end-use efficiency, supporting the use of renewable energy resources as well as sustainable transportation and waste management [5]. One effective approach to avoid CO 2 accumulation into the atmosphere is the recovery of carbon dioxide from flue gases and its recycle by converting to useful chemicals [6, 7]. The Introduction 2 conversion of CO 2 to CO by catalytic hydrogenation has been recognized as a very promising process. In industry, synthesis gas containing H 2 and CO can be used to produce methanol as well as long chain hydrocarbons via the Fischer-Tropsch synthesis. Therefore, the reverse water gas shift (RWGS) reaction (CO 2 + H 2 5 CO + H 2 O) is an important option for CO production. The transformation of CO 2 and H 2 into CO and H 2 O depends upon several factors such as catalyst selection, ratio of CO 2 /H 2 , and reaction temperature and pressure. Therefore, the main focus of this study was to select a suitable catalyst and reaction conditions for CO 2 hydrogenation in a fixed bed reactor. The experiments were designed such that the RWGS reaction could be examined in the forward and reverse direction, i.e. the “normal” water gas shift reaction (CO + H 2 O 5 CO 2 + H 2 ) was also studied. The consecutive CO hydrogenation, i.e. methanation reaction (CO + 3H 2 5 CH 4 + H 2 O) also had to be considered. In heterogeneous catalysis, transport processes (external boundary layer diffusion, pore diffusion) may have an influence on the effective reaction rate, and were also studied in this work in detail. Basic theory and background of the work 3 2. Basic theory and background of the work In the following, some basic aspects of heterogeneous catalysis and phenomena of mass transfer in heterogeneous catalysis (gas-solid diffusion in a fixed bed reactor), deactivation phenomena of catalysts are briefly outlined. Some CO 2 utilization processes including reverse water gas shift, and the thermodynamics of CO 2 conversion are also discussed. 2.1 Heterogeneous catalysis The study of heterogeneous catalysis dates back to the early 1800s where Faraday was one of the first scientists who performed the ability of platinum to facilitate oxidation reactions [8]. After that, the field of heterogeneous catalysis has been grown continuously and attracted several Nobel prizes (for example, in 2007, Gerhard Ertl, a German physical chemist was awarded the Nobel Prize in chemistry for his contribution in the area of surface science). The production of low cost and high quality raw materials, production of transportation fuels, and pollution control are the major areas in which heterogeneous catalysis has shown a remarkable impact. It is well known that the catalyst efficiency is directly proportional to its surface area. Many of the heterogeneous catalysts used in today’s industry consist of one or more catalytically active components deposited on the surface of a support material having a high surface area, high porosity, and a suitable thermal and mechanical strength. Heterogeneous reactions occur in a system in which two or more phases are present (e.g. the solid catalyst and liquids or gases as reactants) and the reactions occur at the interface between these phases. Diffusion processes in heterogeneous catalysis 2.1.1 The chemical reactions of heterogeneous catalysis occur between the adsorbed reactants and the surface of the solid catalyst particle. The diffusion of reactants to the catalyst surface (mostly to the inner surface of pores) and the diffusion of products from this surface are purely physical phenomena. In case of catalytic gas-solid reactions, the rate of reaction within a porous catalyst strongly depends upon the accessibility of reactants to the active sites dispersed throughout the porous structure of the catalyst (Fig. 2-1). Basic theory and background of the work 4 Fig. 2-1: Sequential steps involved in heterogeneous catalytic gas phase reaction. Consider a simple gaseous reaction (A → B) occurring inside a reactor containing porous catalyst particles. In order to convert the reactant A into product B, the following physical and chemical processes are important: 1. Diffusion of reactant A from the bulk gas phase through the boundary layer or stagnant gas film surrounding the catalyst particle to the external surface of the catalyst (film diffusion). 2. Diffusion of reactant A into the porous structure of the catalyst particle to the point where adsorption/reaction takes place (pore or intraparticle diffusion). 3. Adsorption of reactants on the inner surface of the catalyst. 4. Surface reaction of adsorbed reactant A to adsorbed product B at the catalyst surface. 5. Desorption of product B from the inner surface of the catalyst into the pores. 6. Diffusion of formed products through the porous network structure to the external surface of the catalyst (intraparticle diffusion). Basic theory and background of the work 5 7. Diffusion of product B from the external surface of the catalyst through the boundary layer into the bulk gas phase. As a result of these steps, a concentration profile may exist in the catalyst pellet and outside in the film layer. The steps 3, 4, and 5 are chemical processes, which strongly depend on temperature typically with effective activation energies of 20 to 200 kJ/mol. The steps 1, 2, 6, and 7 are diffusional processes or mass transfer with relative low temperature dependence compared to the chemical processes. If these steps are very fast, then there is no resistance for mass transfer from the bulk phase to the external surface area of the particles and to the active sites inside the pore. So the concentration at the internal catalyst surface is the same as that of the bulk phase, and the mass transfer does not affect the reaction rate. But in case that mass transfer is slow relative to the chemical reactions, there may be an influence or control of the overall reaction rate by the mass transfer. The net kinetics obtained from all reaction steps (1 to 7) is thus called effective kinetics or macro-kinetics. Influence of internal and external mass transport on heterogeneously catalysed 2.1.2 reactions In heterogeneous catalytic reactions, the consideration of internal and external mass transfer limitations is very important when experiments are performed to determine intrinsic kinetic parameters. According to Arrhenius law, the intrinsic chemical rate is nearly an exponential function of temperature and the mass transfer rate is less strongly influenced by the temperature, e.g. for gas diffusion the rate is proportional to about T 1.5 . Fig. 2-2 shows an Arrhenius plot of the temperature dependence of the effective rate constant of a catalytic reaction for the three reaction rate controlling regimes. Regime of intrinsic kinetic: At low temperatures, the rate of the chemical reaction is slow compared to mass transport processes. Therefore only a negligible concentration gradient will be established in the exterior and interior of the catalyst particle. The reaction temperature also remains constant over the entire cross-section of the catalyst pellet and also in the boundary layer. This is called as an intrinsic region where the true intrinsic kinetics of reactions is measured. For this region, the slope of the curve ln k m vs 1/T is proportional to the chemical activation Basic theory and background of the work 6 energy (E A ). The reaction rate of gas-solid reactions is often described with accuracy by a power law equation: D E F  4 5    F  E     1  (2.1) The reaction rate usually depends upon the rate constant (k m ) related to the mass of catalyst (m cat ), concentration of reactant (   ), and the intrinsic order of reaction (n). The temperature dependence of the reaction rate constant is given by the Arrhenius equation:  E F  E    A   !" 1  (2.2) In the Arrhenius equation, k m,0 is the frequency factor and E A is the intrinsic activation energy of the chemical reaction. Regime of pore diffusion: At intermediate temperatures, the chemical reaction rate is faster than the inner mass transport (pore diffusion). Between this regime and the regime of intrinsic kinetic, a transition regime is located where the slope of the curve changes with temperature. In the pore diffusion region, the concentration of reactants at the pore mouth is much higher than that inside the pore and drops distinctly. Here the entire catalytic surface is not accessible to the same concentration. Therefore the effective reaction rate will be less as compared to the rate without mass transfer limitations. For this region, the effective activation energy is roughly one half of the true activation energy (E A /2). In general, there is not only the change in effective activation energy but the effective order of the reaction also changes when the transition from kinetic to diffusion control occurs. For an n th order reaction, the effective reaction order approaches the value of (n + 1)/2 in case of strong limitation by pore diffusion. Basic theory and background of the work 7 Fig. 2-2: Typical Arrhenius plot of temperature dependence of the effective rate constant of catalytic reaction; three regimes of reaction rate control for low, intermediate and high reaction temperature. For a first order reaction the effective rate of chemical reaction is given as D E  #$$ F %   E     & 1  (2.3) where  & 'the concentration at the catalyst surface and 6 is the effectiveness factor which is a function of the Thiele modulus. The effectiveness factor 2 is defined as the ratio of the observed rate of reaction to the rate in the absence of any diffusional resistance, i.e. % F (B)*(+  DA(B),-4  D()A ,4)D,4.,B  D()A  -/  B C A,B(+  DA(B),-4 F D E  #$$ D E 1  (2.4) There are several factors which influences the effectiveness factor, such as pore shape and pore structure (micro-macro), particle size distribution, and change in volume upon reaction [9]. The magnitude of the effectiveness factor ranges between 0 and 1 which indicate the relative importance of diffusion and reaction limitations [10]. The internal effectiveness factor varies for different catalyst geometries (see Fig. 2-3) and for different reaction orders (see Fig. 2-4). Basic theory and background of the work 8 Fig. 2-3: Relationship of the effectiveness factor versus generalized Thiele modulus (Eq. (2.7)) for different shapes of catalyst. Fig. 2-4: Relationship of the effectiveness factor versus generalized Thiele modulus (Eq. (2.8)) for simple-order reactions [11]. The effectiveness factor as a function of the Thiele modulus for a flat plate and an isothermal, irreversible first-order reaction is given as: 2324 234 4 2324 234 4 42 422 Effectiveness factor [1] Thiele modulus [2] 567819678A B9CADA EF6AD Basic theory and background of the work 9 % F )(4 C 8 8 1  (2.5) For the spherical particle, an effectiveness factor is given as: % F 0 8 1 0 )(4 C 2 3 8 4 5 0 2 3 8 4 6 1  (2.6) The relation between Thiele modulus and effectiveness factor for a flat plate in Eq. (2.5) gives the approximation that 2 equals 1/8 for large values of the Thiele modulus (8 8 2). The equation of an effectiveness factor for a flat plate can be used in a good approximation for any particle geometry with characteristic length (L p ), which is given by Eq. (2.9). The dimensionless Thiele modulus (8) plays a key role in determining pore diffusion limitations. For general shape, first order and irreversible reaction, the Thiele modulus is given as [9]: 8 F 7 8 9  E  : 8 ;   #$$ 1  (2.7) For a irreversible reaction and order n 7 1, the Thiele modulus is given as: 8 F 7 8 9 1 4 < 0 = 6  E  : 8  > ?  @ ;   #$$ 1  (2.8) where 6 p is the density of particle, D i,eff is the effective diffusivity of species i, and L p is the characteristic length for various shapes which is given as: 7 8 F A-+*A  -/  . C (BA AC)AD4(+  (DA(  -/  . C (BA 1  (2.9) The characteristic length for cylinders is L p = d p /4, for spheres L p = d p /6 and for flat plate L p (where d p is the particle diameter and 2L p is the plate thickness). The point at which a pellet of optimal dimension shows the transition from chemical to diffusion control region is practically important. Small particles have a low characteristic length, which decreases the value of the Thiele modulus (see Eq. (2.7) and Eq. (2.8)) and thus increase the effectiveness factor, and lowers the pore diffusion resistance. On the other hand, a small particle size creates a high pressure drop in a fixed bed reactor. Therefore from Basic theory and background of the work 16 Alloying: In alloying, a combination of two or more metals takes place at high temperature (particularly under reduction conditions). Alloying may change the activity and stability of catalyst. The Cu-Zn alloy formation during the reduction of the catalyst in the temperature range of 230 to 500 °C with a mixture of H 2 and N 2 deactivates the catalyst and reduces its activity for the water gas shift [27]. Other example are formation of RhAl 2 O 4 in Pt-Rh/Al 2 O 3 catalysts in catalytic converters for reduction of car engine emissions, formation of Ni 2 Al 2 O 4 during steam reforming over Ni/Al 2 O 3 , and formation of KAlO 2 in FTS. 2.2 Utilization of CO 2 for production of chemicals and fuels The energy related CO 2 emissions have increased with an average annual growth rate of 1.9% from 1990 to 2007 and the emissions are expected to rise with an annual rate of 1.3% until 2035 [8]. The global CO 2 emissions from fossil fuels combustion increase from near by zero in 1870 to 29.7 billion metric tons in 2007 and are projected to increase up to 42 billion metric tons in 2035 [8]. The analysis of world energy-related carbon dioxide emissions by the consumption of fossil fuels (Fig. 2-6) shows that coal is the largest source of carbon dioxide (12.5 billion metric tons), followed by liquid fuels (11.3 billion metric tons), and natural gas (5.9 billion metric tons). Fig. 2-6: The world energy-related carbon dioxide emissions by fuel type [8]. Basic theory and background of the work 17 To avoid CO 2 emissions, various measures such as recovery, removal, and storage disposal have been proposed. Carbon capture and storage needs large amounts of energy for its capture, transportation, and sequestration. Therefore the utilization of CO 2 in chemical conversion processes may become an important option for sustainable development, mitigation of carbon emissions, and to avoid global warming. The various direct and indirect uses of CO 2 are shown in Fig. 2-7. The potential of CO 2 in the direct use is very low compared to indirect use. The indirect utilization of CO 2 has advantages such as production of value added chemicals, environment friendly processing, and non-hazards process utilization of CO 2 . Fig. 2-7: Various direct and indirect pathways of CO 2 utilization (modified after [5]). Until today, CO 2 is not used in its fullest potential even in the indirect way of utilization because of its high thermodynamic and kinetic stability. The use of efficient catalysts and selective reaction pathways are needed to promote the reaction rate. Undoubtedly, the chemical industry can only make little direct contribution towards the reduction of overall CO 2 emission. According to current estimates, the chemical industry could contribute to convert around 1% of global CO 2 emissions into chemical products [3]. Table 2-2 shows the utilization of CO 2 in various chemical conversion processes in 2006. Basic theory and background of the work 18 The fixation of CO 2 into these organic compounds refers to reactions that use the entire molecule. The detail description for the synthesis of these organic compounds using CO 2 is given in appendix A. All these processes use CO 2 to produce value added products in an environmentally friendly way where utilization potential assisted in terms of less energy use and lower waste production. Table 2-2: Industrial process utilization of CO 2 as a raw material for synthesis of organic compounds [28]. Industrial processes that utilize CO 2 as raw material World capacity per year [million tonnes] Amount of fixed CO 2 [million tonnes] Chemical synthesis: Salicylic acid 0.07 0.025 Urea 143 105 Cyclic carbonates 0.080 0.04 Poly (propylene carbonate) 0.070 0.03 Fuel synthesis: Methanol 20 2 Synthetic natural gas - - Other fuels - - Table 2-2 clearly indicates that the amount of CO 2 utilized today for the production of organic chemicals and fuels (methanol) is very small (around 100 million tonnes) compared to today’s global CO 2 emissions of around 30 billion tonnes. It should also be noted that only around 10% of the global crude oil consumption is used today in the chemical industry. The majorly is used as liquid fuels such as gasoline, diesel and heavy oil. Hence, an effective use of CO 2 with regard to a noticeable reduction of the global net emissions of CO 2 can only be reached, if the CO 2 (e.g. separated from flue gases or in the very far future from air) is utilized for fuels, e.g. by reverse water gas shift and subsequent Fischer Tropsch synthesis. Basic theory and background of the work 19 The reverse water gas shift reaction technology is a simple and effective way to utilize carbon dioxide in many industries [29, 30]. This reaction occurs at high temperatures where CO 2 is converted with H 2 into CO and water. i P   <   j P    k   i   <    j P i       1 ! j l Pmn  o  F  p0 [ =  q O -+ 1  (2.21) Several mostly unwanted parallel and side reactions may also take place: The Sabatier reaction: i P   <   p j P   k   j r   <   = j P i       1 ! j l Pmn  o  F  5 0Y`  q O -+ 1  (2.22) The Bosch reaction: i P   <   = j P   k       <    = j P i       1 ! j l Pmn  o  F  5 \Z [ =  q O -+ 1  (2.23) Boudouard reaction: =s   k       <    i P       1 ! j l Pmn  o  F  5 0t= [ Y  q O -+ 1  (2.24) Methanation reaction: s   <   3 j P   k    j r   <    j P i       1 ! j l Pmn  o  F  5 =ZY  q O -+ 1  (2.25) The reverse water gas shift reaction has been known to chemistry since the mid 1800's but no experimental work was done to reveal its viability [31]. Since the last two decades, studies have been focused on catalytic conversion of CO 2 to industrially important chemicals such as light olefins and liquid hydrocarbons. Depending on the reaction route and catalyst used, it is divided into two groups. One is the hydrogenation of CO 2 to hydrocarbons via methanol synthesis [32, 33] which combines two reaction steps: methanol synthesis from CO 2 and subsequent conversion of methanol to gasoline by the MTG process: i P   <   3 j P   k    j u ij   <    j P i       1 ! j l Pmn  o  F  5 p\ [ `  q O -+ 1  (2.26) = j u ij   k    j u i j u   <    j P i    v     (D-(),B. w [      (2.27) Alternatively, methanol can also be used directly a liquid fuel, but the infrastructure is not yet established. The second route to convert CO 2 into liquid hydrocarbons is Fischer Tropsch synthesis [34-36], which also combines two steps: hydrogenation of carbon dioxide Basic theory and background of the work 20 by reverse water gas shift reaction (Eq. (2.21)) and, then, further hydrogenation of CO to hydrocarbons: i  <  = j P  v  2 5 j P 5 4   <   j P i       1 ! j l Pmn  o  F  5 0`=  q O -+   (2.28) The term (-CH 2 -) represents a methylene group of a paraffin. This route is discussed in some details in the subsequent chapter. 2.3 Concept of production of liquid fuels from CO 2 via reverse water gas shift (RWGS) and Fischer Tropsch synthesis (FTS) The concept for production of liquid fuels from CO 2 via RWGS and FTS was recently described and discussed [37], where some more details can be found. Here only some main aspects should be outlined. Electricity produced by solar energy (or other renewables) can be used to produce liquid fuels like diesel oil with no or little net CO 2 production by the following steps [38-42]: a) Separation of CO 2 from flue gases (or in the long run even from the atmosphere). b) H 2 production by high temperature water (steam) electrolysis and non-fossil electricity: 3 j P s    v    3 j P   <   0 [ ` i P       1 ! j l Pmn  o  F  t3=  q O -+ 1  (2.29) c) CO production by reverse water gas shift (Eq. (2.21)). d) Fischer Tropsch synthesis of hydrocarbons (preferably of diesel oil) (Eq. (2.28)). In summary we get: i P   <   j P s   v    2 5 j P 5 4   <    0 [ ` i P       1 ! j l Pmn  o  F  tpt  q O -+ 1  (2.30) If we compare Eq. (2.30) with natural photosynthesis i P   <   j P s   v    0 Y  x j @P i x   <    i P     1  (2.31) we may regard this process as a technical photosynthesis (TPS, Fig. 2-8). This route is also called carbon capture and conversion (CCC) [42]. Compared to the products of natural photosynthesis (e.g. wood), those of the technical photosynthesis have a much higher energy density related to mass (factor 3) and even more related to volume (factor 6). Basic theory and background of the work 21 Fig. 2-8: The simplified flowsheet of a plant for production of liquid fuels from CO 2 by Fischer Tropsch synthesis and solar energy. The technical photosynthesis may also be helpful for energy storage and transport. Liquid fuels have excellent storage, loading and transport capabilities, and appropriate large-scale storage technologies will be needed in future for the efficient use of renewable energies: (1) Many renewable energy sources (most notably solar and wind) are subject to natural fluctuations. These have to be compensated for by storing excess supply peaks, which increases the efficiency and economic value of renewable energy, and keeps instantaneous electrical generation and consumption in a better balance. (2) The economically most practical method of developing alternative energy sources is to make use of the earth's sunbelt and high-wind zones, but these regions are mostly far away from consumers. Hence, there is a need for a suitable carrier of energy in order to transport the energy from the source to suitable markets. One could also think of using the hydrogen generated by electrolysis directly for storage and transport of solar-electrical energy. But the energy density of liquid hydrogen is only 0.2 toe/m 3 (and even only 0.1 toe/m 3 for compressed H 2 at 700 bar and 25 °C) compared to diesel oil with 0.8 toe/m 3 (toe: tonnes of oil equivalent, 1 toe = 42 GJ). The liquefaction and Basic theory and background of the work 22 transport of H 2 is not easy: High safety requirements and a new infrastructure are needed, e.g. 20% to 30% of the energy content is consumed for liquefaction [43] and 8% for transport as compressed gas via pipelines (per 1000 km) [44]. For a preliminary basic layout of a TPS plant for the production of liquid fuels from CO 2 by solar energy and Fischer Tropsch synthesis (FTS) the following assumptions were made: Hydrogen production H 2 is generated by high temperature solid oxide electrolysis (HTE) at 830 °C. According to Stoots et al. 124 MJ electrical energy are consumed per kg H 2 (= 13.8 MJ per kg converted H 2 O) [45]. If we use the energy released by the exothermic Fischer Tropsch synthesis for the production of saturated steam (22 bar, 215 °C) (see below), the energy needed to overheat the steam (215 5 830 °C) is 1.4 MJ/kg. Hence in total, 137 MJ electrical energy is needed per kg H 2 (= 15.2 MJ per kg converted H 2 O). Desalination of seawater If freshwater is not available, seawater has to be desalinated to produce the feed of the electrolysis. 7 kJ electrical energy is sufficient per kg water, if reverse osmosis is used [46]. This requirement is only 0.04 % of the energy of the subsequent electrolysis. For comparison: 170 kJ/kg is needed, if multi-stage flash distillation is used [44], which is still only 0.9% compared to what the electrolysis consumes. CO 2 production CO 2 has to be separated from flue gases (or on the long run even from air). The inescapable energy requirement (in J/mol CO 2 ) of concentrating CO 2 from a gas mixture is given by the laws of thermodynamics, as we have at least to overcome the difference of the entropy of the mixture and of the pure compounds (see Fig. 2-9): y E F  L  z 3 F 5 { | } ~ 9  L  +4 U { | } ~ W 5 U 0 5 { | } ~ W  9  L  +4 U 0 5 { | } ~ W { | } ~  1  (2.32) Basic theory and background of the work 23 Fig. 2-9: Inescapable energy requirement of concentrating CO 2 from flue gases (11 vol.-% CO 2 ) and air (300 ppmv CO 2 ) at 25 °C (Eq. (2.32)). For separation of CO 2 from flue gases with typically 11 vol-% CO 2 , the minimum energy requirement is 7.8 kJ/mol (177 kJ/kg) and for air 22.4 kJ/mol CO 2 (509 kJ/kg). In the ideal case, each mol of CO 2 is finally converted by FTS into one mol of CH 2 -groups (hydrocarbons) with a heating value of 595 kJ/mol = 42500 kJ/kg. Hence, the energy requirement for CO 2 separation is 3.8% (air) and 1.3% (flue gases) of the energy content of the liquid fuels produced by FTS. Of course in reality, much more energy is consumed at the current status of technology. Here we use a value of 1.2 MJ/kg CO 2 as given by Göttlicher and Pruschek [47] for CO 2 separation from flue gases by chemical absorption. This value is 7 times higher than the inescapable energy requirement. For CO 2 separation from air, a value of 3.6 MJ/kg CO 2 (= 7 E min ) may be taken as a first estimation. In the ideal case of a technical photosynthesis, 3 mol water are converted by electrolysis per mol of CO 2 (Eq. (2.21) and (2.29)), i.e. 1.23 kg H 2 O/kg CO 2 . Hence, for the conversion of 1 kg of CO 2 24.5 MJ are needed for electrolysis of water compared to 1.2 MJ for CO 2 separation from flue gases and 3.6 MJ from air. Reverse water gas shift H 2 and CO 2 are converted to H 2 O and CO by RWGS at a temperature of about 800 °C. At the reactor outlet, the thermodynamic equilibrium is reached as proven by respective experiments with a Ni-catalyst (see section 5.1). The reactor is heated (e.g. electrically). If Basic theory and background of the work 24 heat losses are neglected, the required energy is 41 kJ/mol CO 2 (Eq. (2.21)) or 932 kJ/kg CO 2 . Fischer Tropsch synthesis Production of synthetic fuels via FTS has the potential to produce fuels like gasoline and diesel oil as well as petrochemicals from fossil and renewable sources. In recent years, the availability of cheap natural gas, coal, and biomass has given momentum to FT technology. The worldwide FT plant capacities will increase in future, today with natural gas favoured as feedstock. In 2015, the worldwide annual production of liquid fuels via FT will be around 30 million tonnes, mainly produced in countries like South Africa, Malaysia, and Qatar. Beside of the main reaction of the FT synthesis (Eq. (2.28)), methanation (Eq. (2.25)) is often considered as a separate reaction. The third reaction that plays an important role (at least if iron based catalysts are used) produces unwanted CO 2 by the WGS reaction:   i   <    j P i   k   i P   <   j P       1 ! j l Pmn  o  F  5 p0 [ =  q O -+ 1  (2.33) Two reactor types are currently favoured for FTS, the multi-tubular fixed bed and the slurry bubble column. FT reactors are usually cooled by boiling water. Saturated steam is produced and may be used after further heating as feed for the HTE. A typical FT cooling temperature is 215 °C [48]. The required thermal energy is then 2.7 MJ per kg saturated steam = 48.6 kJ/mol. Hence, 3.1 mol of steam can be generated per mol CO converted by FTS (1jl Pmno = - 152 kJ/mol CO), which is the amount needed for electrolysis. Typical selectivities (in C-%) of the FTS with Fe-catalysts (30% CO conversion per pass, i.e. recycle of unconverted syngas has to be installed) are: 18% CO 2 , 5% methane, 11% C 2 to C 4 , 10% gasoline, 17% diesel and 39% waxes [48-50]. The waxes are further converted (mainly) to diesel oil by mild hydrocracking at 70 bar and 350 °C [51]. All light hydrocarbons (methane to C 4 ) and also the CO 2 formed by the water gas shift reaction in the FTS (Eq. (2.33)) may be recycled to the RWGS unit to be finally converted by steam reforming (reverse of Eq. (2.25)) or by the RWGS (Eq. (2.21)) to CO and H 2 , respectively. In summary we get the following mass balance: For each mol of fresh CO 2 , we have to recycle approximately 0.18 mol CO 2 and 0.18 mol carbon (as C 1 to C 4 ) into the RWGS reactor. Hence, about 2 times more energy is needed to run the RWGS reactor Basic theory and background of the work 25 (1.18 x 41 kJ/mol CO 2 and 0.16 x 206 kJ/mol C as C 1 to C 4 compared to 41 kJ/mol without recycle). Based on the above listed assumptions and facts, an estimation of the energy requirements for the production of liquid fuels from CO 2 by FTS and solar energy is possible (Table 2-3). The energy required for the conversion of CO 2 to CO is calculated by including the conversion of recycled CO 2 and C 1 to C 4 hydrocarbons (counted as CH 4 ), both unwanted byproducts of FTS, and RWGS and steam reforming (reverse of Eq. (2.25)), respectively. Table 2-3: Estimation of energy requirements for the production of liquid fuels from CO 2 by FTS and solar energy (conversion of 1 kg CO 2 to 0.32 kg liquid fuels). Process/energy output Energy [kJ] Comment/assumption Separation of 1 kg CO 2 from flue gas (air) 1200 (3600) Value from [47],which is 7 times the thermodynamic minimum (Fig. 2-9) Conversion of CO 2 to CO (RWGS) 2144 Only counting enthalpy of reaction Desalination of 1.23 kg seawater 9 Reverse osmosis plant HT-electrolysis of 1.23 kg water 18700 13.8 MJ/kg H 2 O electrical energy for electrolysis [45] and 1.4 MJ/kg to overheat saturated steam generated in the FTS (215 5 830 °C) FT synthesis - Heat of reaction is used to generate saturated steam (215 °C, 22 bar) Total required energy 21773 (24173) Production of 0.32 kg liquid fuels 13600 Heating value of fuel = 42500 kJ/kg Process efficiency 62% (= 13600/21773) CO 2 from flue gases 56% (= 13600/24173) CO 2 from air Basic theory and background of the work 32 the production of fuels via Fischer Tropsch synthesis. Exemplarily, the equilibrium composition of syngas is shown in Fig. 2-16. For temperature above 750 °C, atmospheric pressure, and H 2 /CO 2 inlet ratio of 2, the CO 2 hydrogenation can give the required ratio of syngas for the production of fuel. For example, at 800 °C and atmospheric pressure, the equilibrium composition at the reactor outlet is 57% CO, 28% H 2 , and 15 vol.-% CO 2 . Hence, the CO/H 2 ratio is about 2. Fig. 2-16: Equilibrium composition of syngas after removal of water over temperature in CO 2 hydrogenation reaction with methanation (p = 1 atm, H 2 /CO 2 = 2). 2.5 Kinetics of reverse water gas shift reaction and methanation Two general mechanisms have been proposed for methanation of CO 2 over Ni containing catalysts [33, 57-59]. In the first mechanism of methane formation, initially CO 2 undergoes rapid dissociate adsorption with the formation of CO and atomic oxygen on the catalyst surface and, then, further steps coincide with CO hydrogenation steps, where further dissociation of CO occurs into intermediate carbon and oxygen. The intermediate carbon species is then hydrogenated in a chain mechanism to form methane. In this case, CO 2 hydrogenation principle would have to become similar as that of CO hydrogenation. The mechanism proposed for both CO 2 and CO methanation [58] are shown in Table 2-4. Basic theory and background of the work 33 Table 2-4: Mechanism for methane formation from CO 2 and CO [58]. E  1A8C7781A C7!"1 s P  k  s  < s    1#4$ 1  P  k  =    1#$ 1 s  k  s 1#$ 1 s  k    < s  1#$ 1   <   k   1#$ 1   <   k  P  1#$ 1  P  <   k  u  1#$ 1  u  <   k  r  1#$ 1  r  k  r 1#$ 1 i  < j  k ij  1#42$ ij  < j  k j P i  1#44$ j P i  k j P i 1#4$ The second mechanism involves the pathway that does not require initial transformation of CO 2 to CO. Instead, the hydrogenation of CO 2 to methane occurs directly through the formation of intermediates (Table 2-5). This explanation was given on the basis of the difference in specific activity and selectivity with respect to methane formation in CO and CO 2 hydrogenation reaction [59]. According to [59], the steps 13 and 14 are fast and step 15 is slow, and the subsequent steps 16 to 23 are fast and irreversible. Basic theory and background of the work 34 Table 2-5: Mechanism for methane formation without CO formation [59]. E  1A8C7781A C7!"1 s P  k  s P  1 2034  P  k  =    1 20p4 s P  <   k  ss  1 20`4 ss  <   k  s  < s  1 20Y4 s  <   k   < s  1 20t4   <   k  P  1 20K4  P  <   k  u  1 20\4  u  <   k  r  1 2=Z4  r  k  r 1 2=04 ij  < j  k j P i  1 2==4 j P i  k j P i 1 2=34 The direct hydrogenation of CO 2 to hydrocarbons (Eq. (2.39)) using laser generated iron carbide catalysts has been recently investigated by Fiato et al. [60]. This reaction is proposed to proceed via dissociative adsorption of carbon dioxide into CO and oxygen followed by further hydrogenation of the adsorbed species on the catalyst surface via chain reaction mechanism. i P  <  3 j P  v  2 5 j P 5 4 <  = j P i       1 j l Pmn  o  F  5 00Z [ K  q O -+   (2.39) The transformation of CO 2 and H 2 depends upon many factors. Several attempts have been made by several investigators to overcome the limitations of RWGS reaction. Several noble metals (e.g. Pt, Ru, Pd) and transition metals (e.g. Cu, Ni, Fe, Co) supported on different metal oxides (e.g. 8-Al 2 O 3 , CeO 2 , SiO 2 ) have been explored to carry out RWGS reaction [30]. The various catalysts investigated as: Basic theory and background of the work 35 Copper catalyst: A catalyst with 12 wt.-% Cu loaded on alumina has shown a good activity and excellent selectivity to CO in CO 2 hydrogenation [61], where 28% conversion of CO 2 with 100% CO selectivity at 350 °C and atmospheric pressure was established. The reaction was carried out with H 2 /CO 2 feed ratio of 4 and at a space velocity of 100 ml/g cat /min. The exclusive selectivity to CO was obtained on Cu compare to other catalyst such as Ni, Ru, Rh and Pt on alumina. These catalysts produce only CH 4 while Pd catalysts give both CH 4 and CO under the same reaction condition. A catalyst with 5 wt.-% Cu supported on alumina (H 2 /CO 2 feed ratio of 3, space velocity of 100 ml/g cat /min, 220 °C, 3 MPa) produces CO and CH 3 OH with the selectivity of about 87% and 7.5%, respectively [62]. Sometimes, the catalyst support may alter the selectivity of product. With the same amount of Cu on SiO 2 the selectivities are about 87% and 13% for CO and CH 3 OH, respectively while on TiO 2 the values are 92.5% and 4.3% for CO and CH 3 OH, respectively and 2.9% for CH 4 . Similarly, when 5 wt.-% Cu/SiO 2 catalyst operated at 350 °C (pressure of 0.2 bar, gas recycle loop having a H 2 /CO 2 ratio of 4), the selectivity of CO was about 97% with small amount of CH 4 (<1.5%), C 2 H 4 (<0.5%) and C 2 H 6 (<0.4% ) [63]. The addition of a small amount of Ni to a Cu catalyst enhances the catalytic activity but also starts CH 4 formation. When the 5 wt.-% Cu/SiO 2 operated at 280 °C (6 MPa, H 2 /CO 2 ratio of 3, space velocity of 50 ml/g cat /min), the selectivity of CO decreases and the methanol selectivity increases to 76% and 24%, respectively [64]. Iron catalyst: The CO 2 hydrogenation on Fe catalysts has been studied by Lee and coworkers [35, 65-70] and Cubeiro et al. [71]. At 400 °C (2 MPa, space velocity of 31.67 ml/g cat /min), the Fe/Al 2 O 3 catalyst gives about 49% conversion of CO 2 and 12.4% and 87.6% of selectivity towards CO and hydrocarbons, respectively [35]. The addition of potassium to Fe remarkably increases the catalyst activity and selectivity towards C 2+ –hydrocarbons. Under the same reaction condition with K/Fe molar ratio of 0.5, CO 2 hydrogenation activity increased to about 70% while selectivity changes to about 4% and 96% for CO and hydrocarbons, respectively. Further increase in molar ratio of K/Fe to 1 showed a small decrease in activity and hydrocarbon selectivity while the selectivity to CO slightly increased. The addition of alumina itself acts as a promoter when the Fe-Cu-K/Al 2 O 3 Basic theory and background of the work 36 catalyst was prepared by the co-precipitation method [65]. The long term stability test on Fe-K/Al 2 O 3 catalyst showed a significant decrease in CO 2 conversion and selectivity to hydrocarbons [66], where the deactivation was attributed mainly due to the carbonaceous deposition on the catalyst surface. A similar behaviour was obtained with a Fe-Cu-K/Al 2 O 3 catalyst, where the catalyst deactivation was caused by the growth of crystallites and the resulting decrease in dispersion of catalyst components (promoters) [67]. To improve the strength of catalysts, binders are sometimes added. At high temperatures, binders are not chemically inert. The addition of an alumina binder to a Fe-K/Al 2 O 3 catalyst showed excellent activity and selectivity towards higher hydrocarbons (C 5+ ) while activity and selectivity dramatically decreased with silica binder. The reason for this influence on the activity and selectivity was the change in acidity of catalyst, structure and metal-support interaction [68]. The strong metal-support interaction between Fe-K/9-Al 2 O 3 makes the catalyst more active and selective towards long chain hydrocarbons and light olefins [69]. A study on Fe/Al 2 O 3 catalyst preparation methods such as impregnation, precipitation, and physical mixing and catalyst characterization was performed by Cubeiro et al. [71]. The effect of K promoter in both CO 2 and CO hydrogenation was also investigated. Gold catalyst: Gold catalysts supported on various metal oxides such as ZnO, Fe 2 O 3 , CeO 2 , TiO 2 , ZrO 2 La(OH) 3 NiO, and Co 3 O 4 were used by Sakurai et al. [72] for CO 2 hydrogenation. All the catalysts tested for CO 2 hydrogenation were treated at 150 to 400 °C (8 MPa, H 2 /CO 2 ratio of 3, space velocity of 50 ml/g cat /min). Among all the catalysts, Au supported on TiO 2 , Fe 2 O 3 , and ZnO were the most active catalysts for the reverse water gas shift reaction, where the equilibrium conversion of CO 2 was obtained at 400 °C. But even at lower temperatures between 150 and 200 °C, Au/TiO 2 gives equilibrium yield of CO. However, in the temperature range of 150 to 200 °C, Au/Fe 2 O 3 and Au/ZnO catalysts produce much methanol. The activity of the catalyst and the selectivity greatly differ depending upon the reaction pressure and nature of oxide support [73], i.e. when the pressure was decreased from 5 to 0.1 MPa, the CO selectivity increased from 86% to 99% on Au/TiO 2 catalyst (250 °C). Basic theory and background of the work 37 Cobalt catalyst: The CO 2 hydrogenation was also studied over Co Fischer Tropsch catalyst [74, 75]. The Co/SiO 2 catalyst at 210 °C (2.4 MPa, H 2 /CO 2 ratio of 4, space velocity of 83.3 ml/g cat /min) showed deactivation over time and also produces more than 70% of methane [74]. The comparative CO 2 hydrogenation on Fe, Co, and Ni catalyst using Al 2 O 3 as structural support showed highest activity over Co/Al 2 O 3 catalyst. The selectivity order for CO was Fe/Al 2 O 3 > Co/Al 2 O 3 > Ni/Al 2 O 3 [76]. Nickel catalyst: Currently, Ni is one of the most studied catalysts for CO 2 hydrogenation reaction because of its high activity and comparatively low cost. At low temperatures, the CO 2 hydrogenation over Ni catalyst principally produces CH 4 as a main product [58, 77-80]. A kinetic study of CO 2 hydrogenation over Ni/SiO 2 catalyst showed that the activation energy shifts from 89 to 39 kJ/mol as temperature is increased from 227 to 327 °C [58]. The preparation method is also important. The specific activity of Ni/Al 2 O 3 coprecipitated catalyst decreases as the metal loading increases whereas the activity increases with metal loading for an impregnated catalyst [77]. The specific activity (related to Ni only) increases with metal loading up to 30% and then it decreases for Ni/ZrO 2 catalyst prepared by an ultrasound assisted method [78]. During the reaction, transformation of zirconium to zirconium dioxide and the crystallization of metallic Ni particles occurs [80]. At a high Ni content, nickel exists as a bulk NiO and correspondingly produces more methane [81]. The Ni catalyst prepared by coprecipitation method with a loading from 0 to 20% on CeO 2 showed that the 2 wt.-% NiCeO 2 exhibits excellent catalytic activity, selectivity, and stability for the reverse water gas shift reaction [81]. At 600 °C and atmospheric pressure with H 2 /CO 2 ratio of 1 and a quite high space velocity of 2 l/g cat /min, CO 2 conversion obtained over 2 wt.-% Ni/CeO 2 catalyst was about 35% with 100% CO selectivity. The other RWGS reaction catalyst (10 wt.-% NiO/ZnO) showed an even higher activity (CO 2 conversion of 38%, which is very near to equilibrium conversion of 40%) and a selectivity towards CO of about 98% [82]. 38 Objective and scope of the work 39 3. Objective and scope of the work As mentioned in the previous chapter, the consumption of fossil fuels and the release of CO 2 into the atmosphere is continuously growing. Every year billion tons of anthropogenic carbon dioxide are released into the atmosphere. The ability to separate CO 2 from flue gases and to store a several billion tons of CO 2 emitted per year is questionable as are the environmental consequences. Therefore the conversion of CO 2 to useful chemicals and fuels is an attractive option for CO 2 mitigation. A possible avenue for sustainable development is the catalytic conversion of CO 2 to liquid fuels by Fischer Tropsch synthesis and solar or wind energy. The reverse water gas shift reaction (RWGS) is the first step, and the produced CO further converted to liquid fuels by Fischer Tropsch synthesis. The transformation of CO 2 and H 2 depends upon several factors such as catalyst selection, ratio of CO 2 /H 2 , and reaction temperature and pressure. Therefore, the main focus of this work was to use a suitable catalyst and to determine the optimal reaction conditions for CO 2 hydrogenation in a fixed bed reactor. The experiments were designed such that the RWGS reaction could be examined in both the forward and reverse direction. In addition to these experiments, the consecutive reaction of CO to methane was also considered. The studies for all three reactions were performed using a commercial Ni catalyst in a lab scale fixed bed reactor. Since transport processes (external boundary layer diffusion and pore diffusion) may have strong influence, this aspect was also considered. The most available studies of catalytic hydrogenation of CO 2 were performed at low temperatures where methane is the main product. Therefore, in this work CO 2 hydrogenation is studied at high temperatures where the reverse water gas shift reaction is favoured thermodynamically and produces CO and H 2 O as main products. For the basic investigation of catalytic activity, pure alumina and some alumina supported Cr, Zn-Cu, and Ni catalysts were used. After comparing the activity of all these catalysts, the commercial Ni catalyst supplied by Süd-Chemie (spherical pellet, d p = 5 to 7 mm) was chosen for a detailed kinetic study. Objective and scope of the work 40 The parameters that influence the reaction rate such as temperature, reactant concentration, residence time, and particle size of the catalyst were studied in a lab scale fixed bed quartz reactor. To determine the kinetic, the experiments were conducted at conditions sufficiently far away from equilibrium. Not only the RWGS reaction, but also CO hydrogenation (methanation) (CO + 3H 2 5 CH 4 + H 2 O) and WGS reaction (CO + H 2 O 5 CO 2 + H 2 ) were studied in detail. Finally, a one dimensional adiabatic and isothermal fixed bed reactor model was developed at technical conditions (high temperature) over Ni/Al 12 O 19 as well as Al 2 O 3 catalyst system to simulate the performance of hydrogenation of CO 2 to CO for the production of liquid fuels via Fischer Tropsch synthesis. The axial profiles of reactor performance (temperature and conversion) were simulated along the catalytic fixed bed. Experimental method and data analysis 41 4. Experimental method and data analysis In the following, the experimental setup and the procedure used in CO 2 hydrogenation, CO hydrogenation to methane as well as water gas shift reaction, and its corresponding calculation methods are introduced. 4.1 Experimental setup The experimental set-up with reactor apparatus constructed for this study is shown in Fig. 4-1. It mainly consists of a quartz reactor, mass flow controllers, a heating furnace and a temperature control system, a water saturator, a cooling system, and an online gas analyser. Fig. 4-1: Experimental set-up used for CO 2 hydrogenation, CO hydrogenation and water gas shift reaction. A photograph and a schematic diagram of the bench-scale fixed bed reactor (2 cm inner diameter and 45 cm length) used for the experimental studies is shown in Fig. 4-2. In the Experimental method and data analysis 48 methane are the only detectable products of CO 2 hydrogenation reaction, CO 2 conversion is calculated as: i P  B-4AAD.,-4  2  |} ~ 4  F 2 i 4 FJ < 2 j r 4 FJ 2 i 4 FJ < 2 j r 4 FJ < 2 i P 4 FJ  0ZZ 1 (4.3) Similarly, the yield of CO and CH 4 can be calculated (no CO and CH 4 in the feed) as: i  {,A+  2  |} 4  F 2 i 4 FJ 2 i 4 FJ < 2 j r 4 FJ < 2 i P 4 FJ  0ZZ 1 (4.4) j r  {,A+  2  |   4  F 2 j r 4 FJ 2 i 4 FJ < 2 j r 4 FJ < 2 i P 4 FJ  0ZZ 1 (4.5) 4.4.1.2 Determination of the intrinsic kinetic parameters The intrinsic kinetic experiments were carried out in a quartz glass fixed bed reactor with an inner diameter of 2 cm, filled with 1.2 g catalyst of 5 to 7 mm diameter. The isothermal experiments were carried out at atmospheric pressure. The inlet feed passed through the reactor (H 2 /CO 2 = 6) was diluted with 23 vol.-% of nitrogen. The experiments were performed until constant conversion of CO 2 was obtained. For the bimolecular CO 2 hydrogenation reaction (CO 2 + H 2 5 CO + H 2 O), the intrinsic reaction rate over solid catalyst is calculated as: D E  | } ~ F  E  | } ~   |} ~     ~ E 1 (4.6) For the reaction order n 7 1, the reaction rate constant  E|}~ is given as:  E  | } ~  F 2 0 5   ~ 4 @  ? 5 0 2  5 0 4   E   |} ~   @    ~ E 1 (4.7) This calculated  E|}~ in Eq. (4.7) may not be the intrinsic rate constant, which can be confirmed by calculating the effectiveness factor. The modified residence time E as a function of mass of catalyst and volume flow rate at reaction temperature and total pressure of gas is given as:  E F   6 5  2 L   FQ 4 1 (4.8) Experimental method and data analysis 49 The activation energy E A and frequency factor k m,0 can be calculated from the Arrhenius equation:  E  | } ~  F  E  F  A   !" 1 (4.9) 4.4.1.3 Internal mass transport calculations As mentioned in chapter 2, the pore diffusion limitations may occur, and the measured kinetics may not be the intrinsic but the effective one. The effective rate of chemical reaction is calculated with the help of the effectiveness factor 2 is given as: D E  #$$ F %  D E 1 (4.10) The effectiveness factor 2 for n th order irreversible reaction is given by: % F D E  #$$ D E F )(4 C 8 8 1 (4.11) The Thiele modulus for the catalyst with homogeneous distribution and for n th order, irreversible, bi-molecular CO 2 hydrogenation reaction is given as 8 F 7 8  9 1 4 < 0 = 6  E  : 8   |}~   @   ~ E ; |} ~  #$$ 1 (4.12) where the characteristic length L p is the ratio of the volume of the particle to the external surface of the particle (for spherical particle L p = D 8 O3). For a shell catalyst the modified Thiele modulus (8  4 is given as [83] 8  F 7 8  9 1 4 < 0 = 6  E   : 8  |}~   @   ~ E ; |} ~  #$$ 1 (4.13) whereas the characteristic length (L p ) for a shell catalyst is given as 7 8 F  0 5 e 0 5 ) & D 8  h u   D 8 3 1 (4.14) where D 8 is the radius of particle and ) & is the shell thickness of active material. The effective diffusivity of CO 2 in a porous catalyst is given as: Experimental method and data analysis 50 D |} ~  #$$ F E 8 7 8 D |} ~  8FG# F E 8 7 8 e 0 D |} ~  EFQ < 0 D |} ~  IJ h  @ 1 (4.15) In the present study, mass transport in the pores depends both on molecular diffusion and Knudsen diffusion. The molecular diffusion coefficients of CO 2 in the gas mixture were calculated by a computer database program while Knudsen diffusion coefficients were calculated by using Eq. (4.16) as D |}~IJ F   8FG# 3 9 K  9  L M  N |} ~  1 (4.16) where  8FG# is the pore diameter, R is the gas constant, T is the temperature, and N |}~ the molecular weight of CO 2 . The tortuosity factor depends upon nature of interconnecting paths which is influenced by the connectivity, shapes, and specific limiting constrictions. In this work, a typical value for 7 8 of 1.6 was assumed for Ni catalyst. The calculation of the modified Thiele modulus for the shell catalyst needs the modified reaction rate constantU E  W, which is given for a spherical catalyst particle of radius D 8 and shell thickness ) & as:  E  F  8  &   E  6 8 6 &   E 1 (4.17) The factor V p /V s is the ratio of volume of the particle to the volume of the shell: 6 8 6 & F  0 5 e 0 5 ) & D 8  h u   @ 1 (4.18) For the Ni catalyst used in this work V p /V s is about 2.4. By Eq. (4.17) and Eq. (4.13),  E  can be calculated as  E  F  0 5 e 0 5 ) & D 8  h u   @  E 1 (4.19) Insertion of the characteristic length (7 8 ) from Eq. (4.14) and of the modified rate constant ( E  4from Eq. (4.19) into Eq. (4.13), yields the modified Thiele modulus for a shell catalyst having active material only in the shell: Experimental method and data analysis 51 8  F  0 5 e 0 5 ) & D 8  h u   [ ¡  D 8 3  9 1 4 < 0 = 6   E  : 8   |}~   @   ~ E ; |} ~  #$$ 1 (4.20) For the case of) & FD 8 , Eq. (4.19) leads to the same equation that is used for the calculation of the Thiele modulus for a catalyst with a homogeneous distribution of the active component (Eq. (4.12)). 4.4.1.4 Calculation of external mass transport limitations The effective reaction rate due to external mass transfer is calculated as D E  | } ~  #$$ F R  S E  #T U  |} ~ 5  |} ~  & W 1 (4.21) The external mass transfer coefficient 3 (m/s) depends upon the particle size and geometry, molecular diffusion coefficient of gas (D |}~EFQ 4, and the hydrodynamic conditions such as velocity and viscosity of the fluid. 3 as a function of the dimensionless Sherwood number ( ¢ ) for the mass transfer can be calculated as R F 3 C  D |} ~  EFQ  8  1 (4.22) The external surface area A m,ex (m 2 /kg) for spherical particle is S E  #T F Y  8  : 8   1 (4.23) where d p and 6 p are the particle diameter and density of particle, respectively. When gas flows through a catalyst fixed bed, the Sherwood number is given as where 4 b is the bed porosity. The particle Sherwood number ( ¢ £ ) for gas flow around the particle and for laminar flow, correlated to the Schmidt number ( ¢¤ ) and the Reynolds number (Re) is given as [20]: 3 C 8 F = < Z [ Y\  ] 9A  ] 3B  ^ 1 (4.25) The Reynolds number ( ¥ ) depends upon the interstitial velocity u (m/s), the kinematic viscosity of gas ¦ (m 2 /s), and the particle diameter d p (m): 3 C F _ 0 < 0 [ `  2 0 5 H a 4 b  3 C 8  1 (4.24) Experimental method and data analysis 52 9A F *   8 c   1 (4.26) Similarly, Schmidt number ( ¢¤ ) depending upon the kinematic viscosity 7 (m 2 /s) and the molecular diffusion coefficient of gas D |}~EFQ (m 2 /s) as: 3B F c D |} ~  EFQ   1 (4.27) Now the effective reaction rate due to the combined influence of internal and external mass transfer resistances is given in approximation (for an almost irreversible reaction) by d E  #$$ F  e 0 %  f E   |} ~     ~ E < 0 R  g E  #T    ~ h  @ 1 (4.28) CO hydrogenation (methanation) reaction 4.4.2 4.4.2.1 Conversion of CO and yield of CO 2 and CH 4 Similar to CO 2 hydrogenation, the conversion of CO and the yield of CO 2 and CH 4 at given reaction condition is calculated depending upon the outlet carbon percentage values obtained over the analyser. As the methane and carbon dioxide are the main products of CO hydrogenation reactions, the CO conversion is calculated as: i  B-4AAD.,-4  2  |} 4  F 2 i P 4 FJ < 2 j r 4 FJ 2 i 4 FJ < 2 j r 4 FJ < 2 i P 4 FJ  0ZZ 1 (4.29) and the yield of CO 2 and CH 4 in CO hydrogenation is calculated as: i P  {,A+  2  |} ~ 4  F 2 i P 4 FJ 2 i 4 FJ < 2 j r 4 FJ < 2 i P 4 FJ  0ZZ 1 (4.30) j r  {,A+  2  |   4  F 2 j r 4 FJ 2 i 4 FJ < 2 j r 4 FJ < 2 i P 4 FJ  0ZZ 1 (4.31) 4.4.2.2 Kinetic analysis and mass transport calculations The calculations with regard to the intrinsic kinetics and internal and external mass transfer limitations for CO hydrogenation were performed by using the equations applied for CO 2 hydrogenation (see section 4.4.1). Experimental method and data analysis 53 Water gas shift (WGS) reaction 4.4.3 4.4.3.1 Conversion of CO and yield of CO 2 and CH 4 The conversion of CO and the yield of CO 2 and CH 4 in water gas shift reaction were calculated by using the same equations as for the CO hydrogenation (Eq. (4.29) to (4.31)). 4.4.3.2 Kinetic analysis and mass transport calculations The kinetic analysis of the WGS reaction was also done by using the equations of CO 2 hydrogenation reaction (see section 4.4.1). 54 Results and discussion 55 5. Results and discussion As already mentioned in Chapter 4, this work is divided into four parts. In the first part (Chapter 5.1), experimental results are discussed for CO 2 hydrogenation studied by using a fixed bed reactor. This part mainly includes the reaction parameter study, the intrinsic kinetics of CO 2 hydrogenation, and the influence of pore diffusion and film diffusion on the effective reaction rate of CO 2 hydrogenation. The second part (Chapter 5.2) consists of a similar parametric study for CO hydrogenation (methanation reaction). Similar to the first and second part, the third part (Chapter 5.3) includes the parameter study of water gas shift reaction (as the reverse reaction of CO 2 hydrogenation), again the intrinsic kinetics as well as the influence of pore diffusion and film diffusion. The fourth and last part (Chapter 5.4) includes the fixed bed reactor modelling for CO 2 hydrogenation at technical conditions. 5.1 CO 2 hydrogenation (RWGS) The studies on CO 2 hydrogenation (CO 2 + H 2 9 CO + H 2 O) were carried out in a fixed bed quartz reactor over a wide temperature range (300 – 900 °C) and at atmospheric pressure. The commercial Ni catalyst was used. For comparison, 8-Al 2 O 3 was also tested to investigate the catalytic activity of the support only. Effect of reduction temperature on CO 2 hydrogenation (RWGS) 5.1.1 The commercial NiO/Al 12 O 19 (G-90.B) catalyst was reduced at two different temperatures (500 and 800 °C) to examine the effect of the reduction temperature on the catalyst performance and the yield of products. Fig. 5-1 shows the conversion of CO 2 and the yield of CO and CH 4 obtained over the catalyst reduced at 500 and 800 °C. A strong dependence of activity and product yield on the reduction temperature was observed. On the Ni catalyst reduced at a low temperature of 500 °C for 15 h, methane was formed almost exclusively along with CO in the low reaction temperature range (below 600 °C), while on the catalyst samples reduced at high temperature of about 800 °C for 3 h, the methane formation strongly decreased in the same range of reaction temperature. The overall conversion of CO 2 obtained at any reaction temperature is high for the catalyst reduced at 500 °C compared to 800 °C. This could be due to the deactivation of the active species (particularly methane forming species) at high reduction temperatures. Results and discussion 56 Fig. 5-1: CO 2 conversion and CO and CH 4 yield versus temperature; open symbol: catalyst reduced at 800 °C for 3 h; closed symbol: catalyst reduced at 500 °C for 15 h (p = 1 atm, m cat = 1.2 g, catalyst = Ni/Al 12 O 19 , d p = 6 mm, y H2 = 66 vol.-%, y CO2 = 11 vol.-%, rest N 2 , gas flow rate = 138 l/h (NTP)). Over the reduced catalyst both at 500 °C and 800 °C, methanation reaction occurs (most favoured thermodynamically at low temperature) along with the reverse water gas shift (RWGS) reaction. The yield of CO increases continuously with increase in temperature but not the conversion of CO 2 . This is a clear indication of direct methane formation via CO 2 hydrogenation (Sabatier reaction) without intermediate CO formation. The overall CH 4 formation observed could be due to both CO 2 and CO hydrogenation but it is quite difficult to quantify because both reactions occur simultaneously. Fig. 5-2 shows the XRD pattern of the Ni/Al 12 O 19 catalyst (11 wt.-% Ni). The diffractogram (A) for the fresh Ni catalyst shows strong NiO lines at 2B values of 37.3, 43.3, and 62.9 confirming the presence of free nickel oxide [84]. For both the reduced and used catalyst no free nickel oxide lines were observed. On the diffractogram (B) for the reduced catalyst, the lines were observed for nickel crystallites at 44.5 and 51.8 while for the used catalyst (C) a line was observed at 44.5. From the diffractogram (B) it is clear that the catalysts get completely reduced even at low temperature of 500 °C. On the diffractogram (C) for the used catalyst, the free NiO lines were not observed which means that no oxidation of Ni happens by the water formed during the reaction at 500 °C. 0 20 40 60 80 100 300 400 500 600 700 800 900 XCO2, YCO, YCH4 [C-%] Temperature [3C] Equilibrium conversion of CO2 (both to CO and CH4) YCO YCH4 XCO2 Results and discussion 57 Fig. 5-2: XRD patterns for fresh NiO/Al 12 O 19 (A), reduced catalyst at 500 °C (B), and used catalyst at 500 °C (C). To determine the exact temperature needed for the reduction of the Ni catalyst and the number of reducible species present in the catalyst, a TPR experiment was performed by using C HEMBET -3000. TPR analysis begins by passing the analysis gas (10% hydrogen in Results and discussion 64 Ni catalyst 40% CO 2 conversion is reached at about 9 kg s m -3 (Fig. 5-5). Hence, the activity of the Ni catalyst is about 10 times higher. Fig. 5-9: Effect of modified residence time on CO 2 conversion and selectivity of CO and CH 4 (T = 700 °C, p = 1 atm, m cat = 10 g, catalyst = Al 2 O 3 , d p = 1 – 3 mm, y H2 = 66 vol.-%, y CO2 = 11 vol.-%, rest N 2 ). According to the data obtained in this study, methane formation during CO 2 hydrogenation may be due to the direct methanation (CO 2 + 4H 2 5 CH 4 + 2H 2 O) or to the consecutive reaction of CO hydrogenation (CO + 3H 2 5 CH 4 + H 2 O). The consecutive reaction mechanism in CO 2 hydrogenation implies the difficulty that the optimum reaction conditions for each of the reactions are different (high temperature for RWGS and low temperature for CO hydrogenation). Effect of catalyst particle size (RWGS and consecutive methanation) 5.1.5 To determine the effect of particle size on the catalyst performance, particles of Ni catalyst with less than 0.5 mm diameter (crushed catalyst) and 6 mm diameter were used. Fig. 5-10 shows the conversion of CO 2 and the yield of CO and CH 4 at different temperatures. The catalyst particles with diameter of less than 0.5 mm show a higher conversion than the 6 mm particles with the active material only at the shell side which may due to two reasons. One is the fresh catalyst which is not treated at high temperature for long time has higher 2 2 2 2 2 422 2 2 422 42 22 2 XCO2, SCO, SCH4 [C-%] Modified residence time [kg s m-3] Equilibrium conversion of CO2 SCO SCH4 XCO2 Results and discussion 65 activity than the used one (see Fig. 5-4). On the hand the influence of diffusion resistances may be smaller for small particles. This will be analysed in more detail in section 5.1.7. Fig. 5-10: CO 2 conversion and CO and CH 4 yield at different particle size in CO 2 hydrogenation (p = 1 atm, m cat = 1.2 g, catalyst = Ni/Al 12 O 19 , gas flow rate = 138 l/h (NTP), y H2 = 66 vol.-%, y CO2 = 11 vol.-%, rest N 2 ,). Stability of the Ni catalyst for the RWGS reaction at high temperature 5.1.6 The time on stream stability of Ni/Al 12 O 19 catalyst was tested at 900 °C and atmospheric pressure with the total gas flow rate of 138 l/h in the fixed bed reactor. The test was carried out for four days with intermediate cooling cycles. At the beginning of the experiment, the Results and discussion 66 catalyst bed temperature was increased from ambient temperature to 900 °C at the rate of 10 °C/min with a small flow rate of N 2 and then held constant for 30 min to get a stable temperature. Then the required gas composition (CO 2 /H 2 /N 2 ) in the reaction mixture was adjusted and fed to the reactor system. Fig. 5-11 shows the conversion of CO 2 obtained over time on stream of the reaction mixture. At such a high temperature of 900 °C, the CO 2 is practically only converted by the reverse water gas shift reaction and CO is the main product. On the first day of time on stream, the catalyst showed some deactivation where CO 2 conversion decreased from 68% to 66% in 8 h. After that the catalyst was cooled to 300 °C and kept constant overnight. No deactivation was observed on further time on stream, but a small loss in catalyst activity during heating and cooling cycle was observed which may be due to surface modifications. Fig. 5-11: Conversion of CO 2 versus time on stream at 900 °C in CO 2 hydrogenation (p = 1 atm, m cat = 1.2 g, catalyst = Ni/Al 12 O 19 , d p = 6 mm, y CO2 = 11 vol.-%, y H2 = 66 vol.-%, rest N 2 , gas flow rate = 138 l/h (NTP)). Kinetic analysis of the RWGS on the Ni catalyst 5.1.7 Intrinsic kinetics For the kinetic analysis, the concentrations of CO 2 and H 2 were varied (Fig. 5-12 and Fig. 5-13). The temperature was kept low (340 °C) to avoid an influence of mass transport on the effective reaction rate, which then equals the intrinsic rate. Results and discussion 67 Fig. 5-12: Effect of CO 2 concentration on the CO 2 conversion and the CO and CH 4 yield in CO 2 hydrogenation (T = 340 °C, p = 1 atm, m cat = 1.2 g, catalyst = Ni/Al 12 O 19 , d p = 6 mm, y H2 = 66 vol.-%, y CO2 = 11 – 22 vol.-%, N 2 = remaining proportion, gas flow rate = 48 l/h (NTP)). Fig. 5-13: Effect of H 2 concentration on the CO 2 conversion and the CO and CH 4 yield in CO 2 hydrogenation (T = 340 °C, p = 1 atm, m cat = 1.2 g, catalyst = Ni/Al 12 O 19 , d p = 6 mm, y H2 = 44 – 66 vol.-%, y CO2 = 11 vol.-%, N 2 = remaining proportion, gas flow rate = 48 l/h (NTP)). 2    4 4 34 3 3 XCO2, YCO, YCH4 [C-%] CCO2 [mol/m3] 0E 1E 1E2 nCO2= 0 nCO2= 1 2    4 4 3 423 43 XCO2, YCO, YCH4 [C-%] CH2 [mol/m3] 0E 1E 1E2 mH2= 0 mH2= 1 Results and discussion 68 The reaction order with respect to each reactant can be determined by using the following equation:   ~ F 0 5 § f E  |} ~   ¨    ~ E  2  5 0 4   |} ~   @ < 0 © @ @    (5.1)  Fig. 5-14 shows the plot of X CO2 versus C CO2 . The order was calculated by the integral method i.e. the value of the reaction order was determined by the best fit to the measured data. For the experimental conditions used, the reaction order of CO 2 is 0.1. Fig. 5-14: Dependence of CO 2 conversion upon CO 2 concentration (T = 340 °C, p = 1 atm, m cat = 1.2 g, catalyst = Ni/Al 12 O 19 , d p = 6 mm, y CO2 = 11 – 22 vol.-%, y H2 = 66 vol.-%, N 2 = remaining proportion, gas flow rate = 48 l/h (NTP)). Similarly, the reaction order with respect to hydrogen was determined (Fig. 5-15), which leads to a value of 0.4. The reaction rate and the reaction rate constant of CO 2 hydrogenation was calculated by using Eq. (4.6) and Eq. (4.7), respectively. The Arrhenius plot is shown in Fig. 5-16. Only the values of k m,CO2 for T < 410 °C are included, because then the reaction rate is solely determined by the intrinsic kinetics. Above this temperature, the reaction rate is then also controlled by mass transport limitations. 0.02 0.04 0.06 0.08 0.10 0.12 0.14  3  3  3  XCO2 CCO2 [mol/m3] n = -0.1 n = 0.1 n = 0.3 nCO2= 0.1 Fixed point Results and discussion 69 Fig. 5-15: Dependence of CO 2 conversion upon H 2 concentration (T = 340 °C, p = 1 atm, m cat = 1.2 g, catalyst = Ni/Al 12 O 19 , d p = 6 mm, y H2 = 44 – 66 vol.-%, y CO2 = 11 vol.-%, N 2 = remaining proportion, gas flow rate = 48 l/h (NTP)). Fig. 5-16: Arrhenius plot (1/T versus ln k m,CO2 ) for CO 2 conversion (p = 1 atm, m cat = 1.2 g, catalyst = Ni/Al 12 O 19 , d p = 6 mm, y H2 = 66 vol.-%, y CO2 = 11 vol.-%, rest N 2 , gas flow rate = 138 l/h (NTP)). The frequency factor and activation energy (Fig. 5-16) are 5.55 C 10 2 m 1.5 mol 0.5 kg -1 s -1 and 65 kJ/mol, respectively. The value of E A (65 kJ/mol) obtained in this study is within the range found by other authors for CO 2 hydrogenation reaction over Ni catalysts (Table 5-1). 0.07 0.08 0.09 0.10 0.11 0.12   42 44 4 4 4 XCO2 C H2 [mol/m3] 0.4 0.7 mH2= 0.4 Fixed point Results and discussion 70 Table 5-1: Activation energies reported for CO 2 hydrogenation on nickel catalysts. Catalyst Temperature (°C) Pressure (atm) E A (kJ/mol) References 75.3-84.6% Rany Ni 160 – 270 1 88 – 91 [33] 20% Ni/Al 2 O 3 250 – 400 1 87 [76] 3% Ni/SiO 2 227 – 277 1 81 [87] 59% Ni-88 283 – 397 2 – 30 58 – 55 [88] 58% Ni-104 277 – 318 6 – 18 61 [89] 33.6% Ni (G-65) 200 – 230 1 106 [90] 42% NiO/Al 2 O 3 210 – 315 1 80 – 92 [91] 11% Ni (G.90-B) 305 – 410 1 65 This work Influence of internal mass transport (pore diffusion) The pore diffusion limitations were determined for both the shell and non-shell catalyst (homogeneous material). The model parameters required in the calculation of the effective reaction rate are given in Table 5-2. The effective rate of chemical reaction is given as: D E  | } ~  #$$ F %  D E  | } ~ 1 (5.2)  The Thiele modulus for an irreversible, isothermal, nth order reaction and for the spherical particle of 0.5 mm and 6 mm diameter (non-shell catalyst) was calculated by using Eq. (4.12) while for the shell catalyst Eq. (4.17) was used. The effectiveness factor as a function of the Thiele modulus was determined by Eq. (4.11). The Knudsen diffusion coefficient D |}~IJ was calculated by using Eq. (4.16) and the combined pore diffusion coefficient D |}~8FG# was calculated by using Eq. (4.15). The diffusion coefficient of CO 2 in the catalyst pore at 500 °C and 1 bar is shown in Fig. 5-17. For the average pore diameter of 230 nm, the diffusion of CO 2 in the pores of the Ni catalyst is determined by both Knudsen and molecular diffusion (transition area). The Knudsen Results and discussion 71 diffusion dominates the transport phenomenon at pore diameter of less than 100 nm while molecular diffusion dominates above the pore diameter of 10000 nm. Table 5-2: Model parameters used for the calculation of effective reaction rate of CO 2 hydrogenation over Ni catalyst (p = 1 atm). Parameters 2345671 Frequency factor (k m,0 ) 5.55 · 10 2 m 1.5 mol 0.5 kg -1 s -1 Activation energy (E A ) 65 kJ/mol Particle density (6 p ) 1910 kg/m 3 Porosity of particle (4 p ) 0.33 Tortuosity of particle A p (assumption) 1.6 Molecular diffusion coefficient (D CO2,mol ) at 773 K 1.83 · 10 -4 m 2 /s Knudsen diffusion coefficient (D CO2,knu ) at 773 K 4.68 · 10 -5 m 2 /s Pore diffusion coefficient (D CO2,pore ) at 773 K 3.72 · 10 -5 m 2 /s Kinematic viscosity of CO 2 (7 CO2 ) at 773 K 1.63 · 10 -4 m 2 /s Fig. 5-17: Diffusion coefficient of CO 2 in the pores of the Ni/Al 12 O 19 and Al 2 O 3 catalysts determined at 500 °C and 1 bar. Results and discussion 72 The plot of the effectiveness factor versus temperature is shown in Fig. 5-18. Internal mass transfer resistance in the catalyst pores is not negligible for T > 400 °C even for the shell catalyst with 0.5 mm shell thickness of active material (Ni). Fig. 5-18: Effectiveness factor for catalyst particle versus temperature (p = 1 atm, catalyst = Ni/Al 12 O 19 , y CO2 = 11 vol.-%, y H2 = 66 vol.-%, rest N 2 , gas flow rate = 138 l/h (NTP)). Influence of external mass transport (boundary layer diffusion) To determine the influence of film diffusion, the effective reaction rateD E|}~#$$ as a function of external mass transfer coefficient 3 and external surface area per mass of the catalyst A m,ext has been calculated by Eq. (4.21). All the calculations for the external mass transfer coefficient (3), external surface area of the catalyst (A m,ext ), Sherwood number (Sh), Reynolds number (Re), and Schmidt number ( ¢¤ ) are given in section 4.4.1.4. The effective reaction rate due to the combined influence of internal and external mass transport is calculated as: d E  #$$ F  e 0 %  f E  | } ~   |} ~   [ @   ~  [ r < 0 R  g E  #T  U   ~ 5   ~  ª« W h  @ 1 (5.3)  Fig. 5-19 shows a good agreement of the calculations and the measurements for the used shell catalyst. It also reveals that for the hypothetic case of a non-shell catalyst (particle 2324 2342 4322 22 22 22 22 22 22 22 Effectiveness factor (1) Temperature [8C] for particle (dp= 0.5 mm) influence of pore diffusion for catalyst of 0.5 mm shell thickness for particle (dp= 6 mm) homogeneous distribution Results and discussion 73 diameter of 6 mm) the effective reaction rate is lower than for the shell catalyst. For a very small particle size (d p = 0.5 mm) the effective reaction rate is only influenced by pore diffusion at temperature above 700 °C while no influence of external mass transport was observed in the studied range of temperature (not shown here). Considering 6 = 1 in Eq. (5.3) enables to separate the kinetic limitation from internal mass transfer limitation. In Fig. 5-19, the effective reaction rate without internal mass transfer limitation differs significantly from the effective reaction rate with internal mass transport limitation for the shell catalyst as well as for the non-shell catalyst of 6 mm particle diameter. Fig. 5-19 shows that the shell catalyst operates in a mixed regime with both internal and external mass transfer. The catalyst performance is more and more controlled by external mass transfer above a temperature of about 900 °C. Note that Eq. (5.3) is only an approximation because of two reasons: • The reverse reaction (CO + H 2 O 5 CO 2 + H 2 ) is not considered in the pore diffusion term. • For the external mass transfer, the equilibrium concentration is assumed at the external surface. In reality, this value is higher and only reached, if the chemical reaction is very fast. Eq. (5.3) is therefore only exact if only pore diffusion or only external mass transfer plays a role. For the regime in between these two extremes, Eq. (5.3) is only a hopefully good approximation. Furthermore, in Fig. 5-19, intrinsic rate of the chemical reaction (r m ) Eq. (4.6), the reaction rate for diffusion boundary layer (r m,ext ) Eq. (4.21), and the effective reaction rate due to internal and external mass transport (r m,eff ) Eq. (5.3) are shown. The effective reaction rates (r m,eff (calculated) ) were calculated based on the inlet concentrations while the measured values (r m,eff (measured) ) were calculated based on the conversion, inlet concentration, and mass of catalyst. Note that for temperatures between about 500 °C and 900 °C, the equilibrium conversion is not 100% (see Fig. 2-13 for R = 6) and reaches a minimum value of around 75% at 600 °C. This was not considered for the calculations of r m,eff and was only included to calculate r m,ext (FRg E#T 2 ~ 5 ~ª« 4). Results and discussion 80 with temperature until about 450 °C are reached. No catalyst deactivation was observed. In the low temperature range, the reaction is controlled kinetically and no coke formation took place. At low temperatures (below 450 °C) a small amount of CO 2 was produced by the water gas shift reaction [93]. At a very high temperature of about 800 °C, also no catalyst deactivation was observed, but the coke formed in the temperature range of 480 to 710 °C may not completely gasified at 800 °C, and therefore the conversion of CO was low and far from equilibrium. The decrease in conversion due to the catalyst deactivation over time on stream is shown below (see section 5.2.2 (Fig. 5-26)). The coke formation may occur via Boudouard reaction (2CO k C + CO 2 ) which is related to CO dissociation activity in CO hydrogenation [94] and/or direct carbon monoxide hydrogenation (CO + H 2 k C + H 2 O) over the Ni catalyst [95]. Fig. 5-24 shows the equilibrium conversion of CO over the temperature by both Boudouard reaction and direct CO hydrogenation. At high temperatures both reactions are limited thermodynamically. Fig. 5-24: The equilibrium conversion of CO in the Boudouard reaction (CO = 1 mole) and direct hydrogenation reaction (H 2 /CO = 1) as function of temperature and at atmospheric pressure. Stability of the Ni catalyst in CO methanation 5.2.2 Fig. 5-25 shows time on stream behaviour of the Ni catalyst for the CO conversion and the yield of CO 2 and CH 4 (450 °C, atmospheric pressure). The maximum deactivation of the 2 2 2 2 2 422 22 22 22 22 22 22 22 4222 XCO, Equilibrium [C-%] Temperature [3C] 2CO 1 C + CO2 CO + H21 C + H2O Results and discussion 81 catalysts occurs generally in the initial stage of the reaction [93, 96]. In this experiment, no deactivation was observed in 2 h time on stream at 450 °C, and the steady state of reaction was reached within some minutes. When the temperature increased further (keeping all other conditions constant), the coke formation starts above 480 °C, where CO conversion decreases over time on stream in the temperature range of 480 to 710 °C (Fig. 5-26). For example, at 620 °C, the conversion of CO decreases from 7 to 6.3% in 100 min of reaction time. At the temperature of 710 and 810 °C, a small increase in conversion over time on stream was observed which may be due to the gasification of coke formed in the temperature range of 480 to 710°C. Fig. 5-25: CO conversion and the yield of CO 2 and CH 4 versus time on stream in CO hydrogenation (T = 450 °C, p = 1 atm, m cat = 1.2 g, d p = 6 mm, y H2 = 66 vol.-%, y CO = 11 vol.-%, rest N 2 , catalyst = Ni/Al 12 O 19 , gas flow rate = 138 l/h (NTP)). After the completion of the experiment at 810 °C, the reactor temperature was decreased to 300 °C in the presence of small N 2 stream. Then the reactor temperature was increased again from 300 to 700 °C at the rate of 10 °C/min by adding 10% O 2 in N 2 . The presence of CO and CO 2 species in the outlet gas stream during the oxidation of catalyst shows that the coke formed was not completely gasified at 810 °C. 2   4 4 2 2 2 2 2 2 422 42 42 XCO, YCO2, YCH4 [C-%] Time on stream [min] XCO YCO2 YCH4 Results and discussion 82 Fig. 5-26: Time on stream stability of catalyst in CO hydrogenation at different temperature (p = 1 atm, m cat = 1.2 g, catalyst = Ni/Al 12 O 19 , d p = 6 mm, y H2 = 66 vol.-%, y CO = 11 vol.-%, rest N 2 , gas flow rate = 138 l/h (NTP)). Effect of residence time on methanation 5.2.3 The effect of the residence time on the conversion of CO and yield of CO 2 and CH 4 over temperature (300 – 800 °C) was studied at two different gas flow rates of 48 l/h and 138 l/h. In both cases, a similar trend for CO conversion as well as yield of CO 2 and CH 4 was obtained. In the low temperature range of 300 to 450 °C, methane production increases rapidly with temperature and the reaction is controlled kinetically. Above 450 °C, catalyst deactivation due to coke formation may also decrease CO conversion and the yield of CH 4 . At a temperature of about 800 °C and a low gas flow rate of 48 l/h, the conversion of CO almost reached equilibrium. The effect of the residence time on the catalyst performance at constant temperature is given below (section 5.2.4) in detail. Results and discussion 83 Fig. 5-27: CO conversion and CO 2 and CH 4 yield at different gas flow rate versus temperature in CO hydrogenation (p = 1 atm, catalyst = Ni/Al 12 O 19 , d p = 6 mm, m cat = 1.2 g, y H2 = 66 vol.-%, y CO = 11 vol.-%, rest N 2 ). The effect of decrease in CO conversion with increasing temperature is not the result of the equilibrium (at least for T < 700 °C) and also not of a rather slow deactivation by coke formation (sees Fig. 5-26 and Fig. 5-27). Hence, the increase of the temperature obviously leads to a reversible deactivation of the catalyst for the methanation reaction. This strong shift in selectivity with temperature (less methane) was also found in case of CO 2 hydrogenation (Fig. 5-4). It must be noted that the two experiments shown in Fig. 5-27 were done one after the other, i.e. the experiment with the lower flow rate (48 l/h) was done after the experiment with Results and discussion 84 138 l/h. In between these experiments, the catalyst was only oxidised (coke burn-off) and then reduced with H 2 (600 °C). In both experiments, the activity profile is similar. Hence, a high temperature does not lead an irreversible deactivation and only to a strong decrease of the CH 4 selectivity (for T > 400 °C). Whether and to what extent this selectivity shift is induced by the temperature and/or by the formation of coke is still an open question. Effect of residence time at constant temperature 5.2.4 The influence of modified residence time on catalyst performance as well as on the yield of CO 2 and CH 4 was measured by varying the total gas flow rate at 405 °C and atmospheric pressure. As shown in Fig. 5-28, the conversion of CO and the yield of CH 4 and CO 2 increases with an increasing modified residence time. Fig. 5-28: Influence of modified residence time on CO conversion and yield of CO 2 and CH 4 in CO hydrogenation (T = 405 °C, p = 1 atm, m cat = 1.2 g, d p = 6 mm, catalyst = Ni/Al 12 O 19 , y H2 = 66 vol.-%, y CO = 11 vol.-%, rest N 2 ). Interestingly there was no influence of the modified residence time (and thus of the CO conversion) on the selectivity of CO 2 and CH 4 (Fig. 5-29). The extrapolation towards zero CO conversion at zero residence time resulted into no change in selectivity of CH 4 and CO 2 . Hence most probably, direct conversion of CO to CO 2 (2CO + 2H 2 5 CO 2 + H 2 O) may take place (parallel to the CO methanation i.e. CO + 3H 2 5 CH 4 + H 2 O). By comparing the 2  42 4 2  2  2 42 2 2 2 2 2 XCO, YCO2, YCH4 [C-%] Modified residence time [kg s m-3] XCO YCH4 YCO2 Results and discussion 85 conversion data obtained for CO hydrogenation (Fig. 5-28) and for CO 2 hydrogenation (Fig. 5-7) at 405 °C, it is clear that the catalyst has a similar activity for CO and CO 2 hydrogenation, but less CH 4 and more CO are obtained in CO 2 hydrogenation compared to the CH 4 and CO 2 in the CO hydrogenation. Fig. 5-29: CO conversion versus selectivity of CO 2 and CH 4 over Ni catalyst in CO hydrogenation (T = 405 °C, p = 1 atm, 4 m = 10 – 49.7 kg s m -3 , m cat = 1.2 g, d p = 6 mm, catalyst = Ni/Al 12 O 19 , y H2 = 66 vol.-%, y CO = 11 vol.-%, rest N 2 ). Effect of catalyst particle size on methanation 5.2.5 The performance of the Ni catalyst for a particle diameter of less than 0.5 mm and 6 mm over the temperature is shown in Fig. 5-30. For both sizes of catalyst particles a similar trend of CO conversion and yield of CO 2 and CH 4 was obtained. The conversion of CO increases up to 450 °C and then continuously decreases due to the decrease of the methane selectivity by coke formation and/or increasing temperature. 2 2 2 2 2 422 2  42 4 2  2  SCO2, SCH4 [C-%] XCO [C-%] SCH4 SCO2 Results and discussion 86 Fig. 5-30: CO conversion and CO 2 and CH 4 yield at different particle size of catalyst versus temperature in CO hydrogenation (p = 1 atm, m cat = 1.2 g, catalyst = Ni/Al 12 O 19 , y H2 = 66 vol.-%, y CO = 11 vol.-%, rest N 2 , gas flow rate = 138 l/h (NTP)). The previously used catalyst particles of 6 mm diameter in CO 2 hydrogenation study were applied in this experiment. As mentioned in section 5.1.2, the long-time treatment of this catalyst at high temperature reduced its methanation activity. The 0.5 mm size catalyst particles applied in this experiment were used before only in the particle size effect study in CO 2 hydrogenation. Therefore the catalyst particles of 0.5 mm have a higher activity than the 6 mm particles. Due to this reason, in the low temperature kinetic regime a higher conversion of CO and a higher yield of CH 4 was obtained over 0.5 mm size catalyst. Hence, Results and discussion 87 the direct comparison of reaction rate over particle size with different treatment is not possible. Kinetic analysis and influence of internal and external mass transfer on 5.2.6 methanation Intrinsic Kinetics The influence of CO and H 2 concentrations on the methanation is shown in Fig. 5-31 and Fig. 5-32. To determine the reaction order with respect to each reactant, a similar equation that was already applied in CO 2 hydrogenation (Eq. (5.1)) was used:   F 0 5 § f E  |}   ¨    ~   2 £ 5 0 4   |} 8  @ < 0 © @ @  8  (5.4)  The conversion of CO was measured by varying the respective concentration of the reactants CO and H 2 at 340 °C and atmospheric pressure. Fitting this data by Eq. (5.4) shows a negative reaction order for CO (p = -0.3) and a positive order for H 2 (q = 0.7) (for detail see Appendix B.1.1). These values are in good agreement with typical values reported in the literature [97, 98]. Fig. 5-31: Effect of CO concentration on the CO conversion and the CO 2 and CH 4 yield (T = 340 °C, p = 1 atm, m cat = 1.2 g, d p = 6 mm, catalyst = Ni/Al 12 O 19 , y H2 = 66 vol.-%, y CO = 13.2 – 22 vol.-%, N 2 = remaining proportion, gas flow rate = 48 l/h (NTP)). 2 4     3 3 3 XCO, YCO2, YCH4 [C-%] CCO [mol/m3] 0E 1E 1E2 pCO = 1 pCO = 0 Results and discussion 88 Fig. 5-32: Effect of H 2 concentration on the CO conversion and the CO 2 and CH 4 yield (T = 340 °C, p = 1 atm, m cat = 1.2 g, d p = 6 mm, catalyst = Ni/Al 12 O 19 , y H2 = 66 vol.-%, y CO = 11 vol.-%, y H2 = 44 – 66 vol.-%, N 2 = remaining proportion, gas flow rate = 48 l/h (NTP)). The intrinsic rate of CO hydrogenation reaction is therefore given as D E  |} F  E  |}   |}  [ ¬    ~  [ u 1 (5.5)  where k m,CO is the reaction rate constant for CO hydrogenation reaction which can be calculated as: f E  |} F 2 0 5   4 @   5 0 2 £ 5 0 4   E   |} 8  @    ~     Æ C ADA  ¯ B F 5 Z [ 3 ° F Z [ t ± 1 (5.6)  The Arrhenius plot is shown in Fig. 5-33. Only the values of k m,CO for T < 420 °C are included where the reaction rate is exclusively determined by the intrinsic kinetic. The activation energy E A and frequency factor k m,0 calculated from the Arrhenius equation (Eq. (4.9)) are 102 kJ/mol and 2.35 · 10 5 m 1.2 mol 0.6 kg -1 s -1 , respectively. The value of E A obtained in this study (102 kJ/mol) for CO hydrogenation (methanation) is within the range found by other authors using Ni catalysts (see Table 5-5). 2 4       3 423 43 XCO, YCO2, YCH4 [C-%] CH2 [mol/m3] 0E 1E 1E2 qH2 = 1 qH2 = 0 Results and discussion 89 Fig. 5-33: Reaction rate constant (1/T versus ln k m,CO ) of CO hydrogenation (p = 1 atm, m cat = 1.2 g, d p = 6 mm, catalyst = Ni/Al 12 O 19 , y H2 = 66 vol.-%, y CO = 11 vol.-%, rest N 2 , gas flow rate = 138 l/h (NTP)). Table 5-5: Activation energies reported in literature for CO hydrogenation on Ni catalysts. Catalyst Temperature (°C) Pressure (atm) E A (kJ/mol) References Monolithic Ni 200 – 350 6.8 104 – 89 [98] 3% Ni/SiO 2 227 – 277 1.38 96 [99] Ni/MgAl 2 O 4 205 – 290 1.4 96.7 [100] 10% Ni/SiO 2 270 – 320 1 112 [101] 1.04Ni/K-Al 2 O 4 339 – 364 1 94  ²  3 [102] 11% Ni (G.90-B) 305 – 345 1 102 This work Influence of mass transfer limitation The influence of internal mass transport on the effective reaction rate of CO hydrogenation was determined for the shell catalyst as well as non-shell catalyst (6 mm particle diameter with homogeneous distribution). For external mass transport, external surface per mass of 2242 (42 ( ( ( ( ( 23224 23224 23224 232244 23224 23224 T [8C] ln km,CO [m1.2 mol-0.6 kg-1 s-1] 1/T [1/K] Trendline intrinsic reaction rate constant EA= 102 kJ/mol Km,o = 2.35 2 105m1.2 mol0.6 kg-1 s-1 Results and discussion 96 Fig. 5-38: Influence of modified residence time on CO conversion and the yield of CO 2 and CH 4 in water gas shift reaction (T = 405 °C, p = 1 atm, m cat = 1.2 g, d p = 6 mm, catalyst = Ni/Al 12 O 19 , y H2O = 44 vol.-%, y CO = 11 vol.-%, rest N 2 ). Stability of the Ni catalyst in the WGS reaction 5.3.3 5.3.3.1 Low temperature stability In the water gas shift reaction, the Ni catalyst behaves differently at different temperature. No deactivation was observed below 450 °C, but the catalyst shows deactivation due to coke formation in the temperature range of 450 to 700 °C. So, depending upon this behaviour the time on stream stability of catalyst was studied at both low and high temperatures. The low temperature catalytic stability of the Ni/Al 12 O 19 catalyst in the WGS reaction was tested for about 100 min at 340 °C and 450 °C, and atmospheric pressure. Fig. 5-39 shows the conversion of CO over the time on stream. The Ni/Al 12 O 19 catalyst showed a good stability at these low temperatures (340 °C and 450 °C). 0 10 20 30 40 0 10 20 30 40 XCO, YCO2, YCH4 [C-%] Modified residence time [kg4s/m3] T = 405 0C XCO 3 YCO2 YCH4 Results and discussion 97 Fig. 5-39: CO conversion versus time on stream in WGS reaction at low temperature (p = 1 atm, m cat = 1.2 g, catalyst = Ni/Al 12 O 19 , d p = 6 mm, y CO = 11 vol.-%, y H2O = 44 vol.-%, rest N 2 , gas flow rate = 48 l/h (NTP)). 5.3.3.2 High temperature stability The time on stream behaviour of the Ni catalyst in the high temperature range of 500 °C to 900 °C and at atmospheric pressure for water gas shift reaction is shown in Fig. 5-40. It was difficult to determine the amount of coke formed via thermogravimetric analysis because a very small amount of coke is formed and the mass loss by burning of the coke and the mass gain by catalyst oxidation takes place simultaneously. The experimental procedure was as follows: At first, the water gas shift reaction was performed at a certain temperature, e.g. at 500 °C (Fig. 5-40). Then the temperature was increased to 900 °C, where not only the WGS reaction takes place but also the gasification of the coke formed at the low temperature experiment (Fig. 5-41). Fig. 5-40 indicates that at 900 °C, no coke is formed, i.e. the coke that may be formed is gasified with steam at an appropriate rate. Hence, the catalysts used for WGS reaction at lower temperatures for about 100 min regain their activity by gasification, if the temperature of 900 °C is adjusted afterwards. This is shown in Fig. 5-41. 2  4   2 2 2 2 2 2 422 42 XCO [C-%] Time on stream [min] 450 4C 340 4C Results and discussion 98 Fig. 5-40: CO conversion versus time on stream in WGS reaction at high temperatures (p = 1 atm, m cat = 1.2 g, catalyst = Ni/Al 12 O 19 , d p = 6 mm, y CO = 11 vol.-%, y H2O = 44 vol.-%, rest N 2 , gas flow rate = 48 l/h (NTP)). Fig. 5-41: CO conversion versus time on stream in WGS reaction at 900 °C after the lower temperature experiments (p = 1 atm, m cat = 1.2 g, d p = 6 mm, y CO = 11 vol.-%, y H2O = 44 vol.-%, rest N 2 , catalyst = Ni/Al 12 O 19 , gas flow rate = 48 l/h (NTP)). In another experiment, the freshly reduced Ni catalyst (800 °C) was applied for 6 h in the reaction at 650 °C where the catalyst shows deactivation with time on stream (Fig. 5-42). Then this used catalyst was oxidised in the same reactor using 5.5 vol.-% O 2 in N 2 by increasing the temperature from 300 to 700 °C at a rate of 10 °C/min. The presence of CO 2 4 2  2  2 2 2 2 2 422 42 42 XCO [C-%] Time on stream [min] 500 4C 700 4C 800 4C 600 4C900 4C 2 42 2 2 2 2 2 2 2 2 2 2 422 42 XCO [C-%] Time on stream [min] 2215E 2215E 2215E 2215E TWGS = 900 4C Results and discussion 99 and CO 2 species in the outlet gas stream during oxidation of catalyst (coke burning) were analysed using a micro-gas analyser and the amount of coke formed was quantified from the CO 2 produced during the catalyst oxidation. The activity of the catalyst was then recovered by oxidation/reduction. The regenerated catalyst was again used for 3 h and the same procedure was repeated to confirm the coke formation. Fig. 5-42: Time on stream stability and coke formation over freshly reduced catalyst in WGS reaction (T = 650 °C, p = 1 atm, m cat = 1.2 g, d p = 6 mm, y CO = 11 vol.-%, y H2O = 44 vol.-%, rest N 2 , catalyst = Ni/Al 12 O 19 , gas flow rate = 48 l/h (NTP)). Fig. 5-42 clearly shows that the regeneration by coke burn-off is possible. The amount of coke is rather small compared to the carbon (as CO) that has passed through the reactor. During the 6 h experiment, about 17 g of carbon (as CO) have entered the reactor but only 0.45 mg of coke was formed. Nevertheless this small amount of coke is sufficient for a strong deactivation of the catalyst, although the amount of coke (0.04 mmol C) is still small compared to the amount of Ni (2 mmol), but at least in the same order of magnitude (and not all of the Ni material is accessible by reactants). Results and discussion 100 Kinetic analysis and influence of internal and external mass transfer on the WGS 5.3.4 reaction Intrinsic kinetics The reaction orders of the water gas shift reaction with respect to reactant and product components over the Ni/Al 12 O 19 catalyst were determined by fitting the experimental data to Eq. (5.8).   F 0 5 § f E  |}   ¨    ~ } &    ~ J  |} ~ ³  |} G  @  2 d 5 0 4 < 0 © @ @  G  (5.8)  The respective experimental results are shown in the Fig. 5-43 to Fig. 5-46. The reaction orders with respect to reactant CO and H 2 O are 0.8 and 0.4 while the products H 2 and CO 2 shows a -0.15 and -0.1 order dependencies on the reaction rate of WGS reaction. (see Appendix B.1.2). Both of these negative order dependencies of products were not considered in the intrinsic rate calculation. Fig. 5-43: Effect of CO concentration on the CO conversion and the CO 2 and CH 4 yield in WGS reaction (T = 340 °C, p = 1 atm, m cat = 1.2 g, d p = 6 mm, y H2O = 44 vol.-%, y CO = 5.5 – 14.3 vol.-%, N 2 = rest proportion, catalyst = Ni/Al 12 O 19 , gas flow rate = 48 l/h (NTP)). 2  42 4 2  2 432 434 34 34 XCO, YCO2, YCH4 [C-%] C CO [mol/m3] 0E 1E 1E2 rCO = 1 rCO = 0 Results and discussion 101 Fig. 5-44: Effect of H 2 O concentration on the catalyst activity and the CO 2 and CH 4 yield in WGS reaction (T = 340 °C, p = 1 atm, m cat = 1.2 g, d p = 6 mm, y CO = 11 vol.-%, y H2O = 34 – 54 vol.-%, N 2 = rest proportion, catalyst = Ni/Al 12 O 19 , gas flow rate = 48 l/h (NTP)). Fig. 5-45: Effect of H 2 addition into the feed gas of WGS reaction (T = 340 °C, p = 1 atm, m cat = 1.2 g, d p = 6 mm, catalyst = Ni/Al 12 O 19 , y CO = 11 vol.-%, y H2O = 44 vol.-%, y H2 = 0 – 18.75 vol.-%, N 2 = rest proportion, gas flow rate = 48 l/h (NTP)). 2  42 4 2  2 3 3 4232 XCO, YCO2, YCH4 [C-%] CH2O [mol/m3] 0E 1E 1E2 sH 2 O = 0 sH2O = 1 2    4 4 4 4  2 43 3 3 XCO, YCO2, YCH4 [C-%] C H2 [mol/m3] 0E 1E 1E2 uH2= 0 Results and discussion 102 Fig. 5-46: Effect of CO 2 addition into the feed gas of WGS reaction (T = 340 °C, p = 1 atm, m cat = 1.2 g, d p = 6 mm, catalyst = Ni/Al 12 O 19 , y CO = 11 vol.-%, y H2O = 44 vol.-%, y CO2 = 0 – 16.67 vol.-%, N 2 = rest proportion, gas flow rate = 48 l/h (NTP)). The intrinsic reaction rate of WGS reaction has been calculated as D E  |} F  E  |}   |}  [ n    ~ }  [ r F  E    A    !"   |}  [ n    ~ }  [ r 1 (5.9)  where k m,CO is the reaction rate constant for WGS reaction which can be calculated as: f E  |} F 2 0 5   4 @   [ n 5 0 2 Z [ K 5 0 4   E   |}  [ n  @    ~ }  [ r 1 (5.10)  The temperature dependency of k m,CO is shown in Fig. 5-47. To determine the intrinsic kinetic, only the values of k m,CO for T < 345 °C are included where no mass transport or thermodynamic limitations occur. The frequency factor k m,0 and apparent activation energy of reaction E A are 3.46 · 10 5 m 3.6 mol -0.2 kg -1 s -1 and 96 kJ/mol, respectively. The activation energy data reported in literature for the water gas shift reaction over Ni containing catalyst is given in Table 5-7. The value of E A (96 kJ/mol) obtained in this study is within the range found by other authors. 2    4 4 4 4  2 43 3 3 XCO, YCO2, YCH4 [C-%] C CO2 [mol/m3] 0E 1E 1E2 vCO2= 0 Results and discussion 103 Fig. 5-47: Reaction rate constant (1/T versus ln k m,CO ) of the WGS reaction (p = 1 atm, m cat = 1.2 g, d p = 6 mm, catalyst = Ni/Al 12 O 19 , y H2O = 44 vol.-%, y CO = 11 vol.-%, rest N 2 , gas flow rate = 48 l/h (NTP)). Table 5-7: Activation energies reported for water gas shift reaction on nickel catalysts. Catalyst Temp (°C) Pressure (atm) E A (kJ/mol) Ref. 5% Ni/Al 2 O 3 monolith 300 – 1000 1 85 [104] 5% Ni/5% Ce/Al 2 O 3 monolith 300 – 1000 1 85 [104] 5 at.% Ni/Ce(10% La)O x 275 – 300 1 38 [105] Black NiO 200 – 300 1 96 ²  p [106] Green NiO 200 – 300 1 89 ²  p [106] 11% Ni (G.90-B) 305 – 345 1 96 This work Influence of mass transport To determine the influence of mass transport on the effective reaction rate of water gas shift reaction, all the calculations were made by using the equations from section 4.4.1. Similar to 22422222 ( (3 ( (3 ( (3 ( (3 ( 232242 23224 23224 23224 23224 23224 T [8C] ln km,CO [m3.6 mol-0.2 kg-1 s-1] 1/T [1/K] EA= 96 kJ/mol km,0 = 3.46 61105m3.6 mol-0.2 kg-1 s-1 Trendline intrinsic reaction rate constant Results and discussion 104 CO 2 and CO hydrogenation, the influence of mass transfer in WGS reaction was calculated for both shell catalyst (thickness of 0.5 mm) and non-shell catalyst (d p = 6 mm). The effective reaction rate due to the influence of both internal and external mass transport was calculated as: d E  #$$ F  e 0 %  f E  |}   |}   [ n   ~ }  [ r < 0 R  g E  #T  U   5    ª« W h  @ 1 (5.11)  The model parameters required in the calculation of effective reaction rates of water gas shift reaction are given in Table 5-8. The detailed pore diffusion calculation for Ni/Al 12 O 19 catalyst is given in appendix B.2 (Fig. B-8). Table 5-8: Model parameters used for the calculation of effective reaction rate of the water gas shift reaction over Ni catalyst (p = 1 atm). Parameters Values Frequency factor (k m,0 ) 3.46 · 10 5 m 3.6 mol -0.2 kg -1 s -1 Activation energy (E A ) 96 kJ/mol Particle density (6 p ) 1910 kg/m 3 Porosity of particle (4 p ) 0.33 Tortuosity of particle A p (assumption) 1.65 Molecular diffusion coefficient of CO (D CO,mol ) at 773 K 1.07 · 10 -4 m 2 /s Knudsen diffusion coefficient (D CO,knu ) at 773 K 5.86 · 10 -5 m 2 /s Pore diffusion coefficient (D CO,pore ) at 773 K 3.78 · 10 -5 m 2 /s Kinematic viscosity of CO (7 CO ) at 773 K 8.17 · 10 -5 m 2 /s In the water gas shift reaction, the influence of pore diffusion on the effective reaction rate was evaluated from effectiveness factor 6. Fig. 5-48 shows the effectiveness factor versus temperature for the shell as well as for a hypothetic non-shell catalyst. For both catalysts the effectiveness factor decreases strongly with increasing temperature close to 345 °C. Fig. 5-49 compares the effective reaction rates (considering 6 = 1 and with actual 6). There is a clear influence of pore diffusion for both the shell and non-shell catalyst. Results and discussion 105 Fig. 5-48: Effectiveness factor for catalyst versus temperature (p = 1 atm, m cat = 1.2 g, catalyst = Ni/Al 12 O 19 , y CO = 11 vol.-%, y H2O = 44 vol.-%, rest N 2 , gas flow rate = 48 l/h (NTP)). Fig. 5-49: Influence of internal and external mass transport on the effective reaction rate of WGS reaction (p = 1 atm, m cat = 1.2 g, y CO = 11 vol.-%, y H2O = 44 vol.-%, rest N 2 , catalyst = Ni/Al 12 O 19 , gas flow rate = 48 l/h (NTP)). 23224 2324 234 4 22 22 22 22 22 22 Effectiveness factor (1) Temperature [8C] pore diffusion region *D197D8 6A1#911""$ C"'AA3!1!8D438 *D197D8 6A1*1231"" !CA6618C &A!!11 Results and discussion 112 Fig. 5-51: Axial profiles of CO 2 conversion for the Ni/Al 12 O 19 catalyst and three different cases (Eq. (5.19);  E|}~ v» Rv», and regular model (H 2 /CO 2 = 3). The simulated adiabatic temperature profiles at different gas superficial velocities along the axial reactor coordinate are shown in Fig. 5-52 both for Ni/Al 12 O 19 and Al 2 O 3 as catalyst. The reaction temperature decreases along the length due to the prevailing endothermic Results and discussion 113 RWGS reaction, i.e. the energy consumption by the reaction leads to a cooling of the gas. For an inlet gas temperature of 1200 °C and the highest superficial gas velocity of 15 m/s (at T 0 ), the adiabatic final temperature of 998 °C is reached at a reactor length of about 0.4 m for the Ni catalyst and at about 25 m for the Al 2 O 3 catalyst. The respective residence times (related to the empty tube) are only about 30 ms for the Ni catalyst and about 2 s for the Al 2 O 3 . For lower gas velocities, the required lengths (and residence times) are smaller. Fig. 5-52: Axial temperature profiles of an adiabatic reactor for Ni/Al 12 O 19 and Al 2 O 3 as catalysts (H 2 /CO 2 = 3). Similarly, the axial profiles of the CO 2 conversion are shown in Fig. 5-53 for the simulated adiabatic and isothermal case both for the Ni and Al 2 O 3 catalyst. The conversion of CO 2 and Results and discussion 114 thus the yield of CO increases along the reactor length and reaches the maximum (equilibrium) value at the reactor outlet. As the gas superficial velocity increases, the reactor length required to reach the equilibrium also increases. As already stated, an increasing gas velocity also increases the required length of the reactor to reach (almost) the final equilibrium value. For the Ni catalyst and an inlet temperature of 1200 °C, an increase from 0.5 m/s to 15 m/s leads to an increase of the reactor length from about 0.013 m to 0.4 m for the adiabatic reactor and from 0.01 m to 0.3 m for the isothermal reactor. Fig. 5-53: Conversion of CO 2 over Ni/Al 12 O 19 and Al 2 O 3 catalysts at different gas superficial velocities versus axial reactor co-ordinate (H 2 /CO 2 = 3). The conversion of CO 2 to CO obtained at the reactor outlet is about 82% for the adiabatic reactor and 87% for the isothermal reactor. Results and discussion 115 Fig. 5-54 shows the conversion of CO 2 along the reactor length at different H 2 /CO 2 ratios for the adiabatic and isothermal simulation (again both for the Ni and Al 2 O 3 catalyst). As the ratio H 2 /CO 2 increases, the conversion of CO 2 raises along the reactor length. For example, if the inlet ratio of H 2 /CO 2 is increased from 1 to 6, the conversion of CO 2 at the reactor outlet increases from about 55% to 91% for the adiabatic reactor and from 62% to 93% for the isothermal reactor. But high inlet H 2 /CO 2 ratios also increase the H 2 /CO ratio of the produced syngas, which then may not be optimal for a subsequent Fischer Tropsch synthesis. The H 2 /CO ratio of the syngas at the reactor outlet (R syngas ) for different H 2 /CO 2 ratios of the feed to the RWGS reactor is shown in Fig. 5-55. Fig. 5-54: Conversion of CO 2 over Ni/Al 12 O 19 and Al 2 O 3 catalysts at different inlet gas composition versus axial reactor co-ordinate (u s = 15 m/s, R = H 2 /CO 2 ). Results and discussion 116 Fig. 5-55: Molar H 2 /CO ratio of the syngas (R syngas ) at the reactor outlet for different molar H 2 /CO 2 ratios of the feed (R feed ) to the RWGS reactor (p = 1 bar). Fig. 5-55 indicates that an almost optimal H 2 /CO ratio of the syngas of slightly above two (to account for the methanation, which always takes place during Fischer Tropsch) is reached for a H 2 /CO 2 ratio of the feed of about three. This ratio was therefore used for further simulations. As mentioned before, the reactor length required to get a particular conversion depends upon the gas inlet temperature, the gas superficial velocity, and the inlet gas composition. For the RWGS reaction over the Ni/Al 12 O 19 catalyst (inlet temperature of 1200 °C, inlet H 2 /CO 2 ratio of 3), the reactor length required to reach 99% of the equilibrium conversion in the adiabatic and isothermal case is shown in Fig. 5-56. Fig. 5-57 shows the similar pattern for Al 2 O 3 catalyst. As the effective rate of the RWGS reaction over Al 2 O 3 is by about a factor of 40 lower compared to the Ni catalyst, the reactor length required is much higher compared to Ni/Al 12 O 19 catalyst to get the same conversion. 2 4       4 Rsyngas = H2/CO (reactor outlet) Rfeed = H2/CO2 Isothermal Adiabatic T0= 1200 4C T0= 1000 4C Results and discussion 117 Fig. 5-56: Reactor length required for 99% of equilibrium conversion in adiabatic reactor (X CO2 = 81%) and isothermal reactor (X CO2 = 86%) versus superficial gas velocity over Ni/Al 12 O 19 catalyst (H 2 /CO 2 = 3, p = 1 bar). Fig. 5-57: Reactor length required for 99% of equilibrium conversion in adiabatic reactor (X CO2 = 81%) and isothermal reactor (X CO2 = 86%) versus superficial gas velocity over Al 2 O 3 catalyst (H 2 /CO 2 = 3, p = 1 bar). Results and discussion 118 Pressure drop in fixed bed technical RWGS reactors 5.4.3 According to Ergun [107], the pressure loss in a fixed (packed) bed is given by z B a F e / a  7  8  : $  * & P = h 1 (5.27)  where : $ is the density of the fluid (kg/m 3 ), / a the friction factor of a packed bed, which is given by the following equation based on the particle Reynolds number, 9AF* &  8 Oc: / a F 2 0 5 ¼ a 4 ¼ a u 1 3 [ ` < 2 0 5 ¼ a 4 3ZZ 9A 6 1 (5.28)  For non-spherical particles an equivalent particle diameter is used  8 F e  Y  6 8 S E  #T h 1 (5.29)  which is the diameter of a sphere with the same external surface area per unit volume as the actual particle. For a packed bed of spheres (equal diameter, porosity ¼ a = 0.4), Eq. (5.28) simplifies to / a F 33 < 0tZZ 9A 1 (5.30)  The particle diameter and the fluid velocity have a strong influence onzB a . As we can see by insertion of Eq. (5.30) into Eq. (5.27), 1B a is proportional to * & O 8 P for low values of 9Aand to * &P O 8 for high values. The production of syngas by CO 2 hydrogenation would most probably been done at the typical pressure of the subsequent production of liquid fuels via Fischer Tropsch synthesis. Hence, a total pressure of 30 bar, a maximum pressure drop of 1 bar, and an inlet temperature of 1200 °C are assumed for the final simulation of a technical RWGS reactor. In addition, only adiabatic operation is subsequently considered, as this type of operation is much easier to realize in a technical reactor compared to isothermal operation. The fluid density and viscosity are calculated at 1100 °C (average temperature in adiabatic reactor). Values of the required parameters to calculate zB a are given in Table 5-11. Fig. 5-58 shows that the pressure drop per meter length strongly increases with increasing superficial gas velocity. In case of the Al 2 O 3 catalyst and an adiabatic fixed bed reactor, the superficial gas Results and discussion 119 velocity should be less than 1.7 m/s to limit the pressure drop to the assumed value of 1 bar. The reactor length is then 2.8 m (to get 99% of the equilibrium conversion) (Fig. 5-57). At this velocity, zB a O7 is about 0.38 bar/m and so 1B a D 1 bar over the total reactor length. In case of the Ni/Al 12 O 19 catalyst and at same superficial gas velocity of 1.7 m/s, a reactor length of 0.05 m (Fig. 5-56) would be sufficient and zB a would only be about 0.02 bar. The respective gas velocity for the Ni/Al 12 O 19 catalyst, where the assumed limiting value of zB a of 1 bar is reached, is 6.7 m/s (reactor length of 0.18 m). Table 5-11: Model parameters used to calculate the pressure drop of a RWGS reactor. Parameters (30 bar, 1100 °C) Values Kinematic viscosity (7), m 2 /s 6 · 10 -6 Fluid density (3 f ), kg/m 3 4.65 Particle diameter (d p ), m 0.006 Particle Reynolds number (Re) at u s = 1 m/s 1000 Friction factor (f b ) at u s = 1 m/s 34.7 Fig. 5-58: Pressure drop in fixed bed reactor per meter length for different superficial gas velocities. 23224 2324 234 4 42 422 234 4 42 5pb/L [bar/m] Superficial gas velocity, us[m/s] T = 1100 4C p = 30 bar 5pb 31 bar (Al2O3) 5pb 30.02 bar (Ni/Al12O19) 5pb 31 bar (Ni/Al12O19) Results and discussion 120 Estimation of the size of the RWGS reactor with regard to a subsequent Fischer 5.4.4 Tropsch fixed bed reactor It is interesting to estimate the size of an adiabatic RWGS reactor (both for the investigated Niand the Al 2 O 3 catalyst) needed for a subsequent Fischer-Tropsch (FT) reactor. According to the data given by Jess and Kern [48], the parameters of a technical FT reactor are as follows: o The FT reactor is a multitubular reactor with 8000 tubes each with a diameter of 5 cm. o The total volume rate of fresh (dry) syngas is about 160,000 m 3 /h (NTP). o The total pressure is around 30 bar. The reactor length required to get 99% of the equilibrium conversion of CO 2 is calculated as 7 mm  ½|} P  # F 7   * & 111111111111(i: Nior Al 2 O 3 -catalyst)1 (5.31)  And the length to reach a certain zB a (in this calculation zB a = 1 bar) as 7 z 8 ¾ F =  z B a   8 / a  : $   * & P 1 (5.32)  For the estimation of the design of the adiabatic RWGS reactor, the following assumptions were made: o For the assumed volume rate of dry syngas of 160,000 m 3 /h (NTP), the volume rate through the RWGS reactor is 201,000 m 3 /h (NTP), because for a CO 2 conversion of 81% (reached in adiabatic operation) 20% of the syngas is water, which is separated. At a mean temperature of 1100 °C and 30 bar, this corresponds to 34,000 m 3 /h (= 9.4 m 3 /s). o For the Ni/Al 12 O 19 catalyst, the residence time required to get 99% of the equilibrium conversion of CO 2 is 0.027 s (see Fig. 5-56); the respective value for the Al 2 O 3 catalyst is 1.65 s (see Fig. 5-57). (Remark: Pore diffusion is only determined by Knudsen diffusion. Hence, there is practically no influence of the total pressure on the effective reaction rate (constant) and thus also not on the required residence time.) o For the pressure drop in the RWGS reactor a limiting value of 1 bar is assumed. The dimensions of the fixed bed of the RWGS reactor are then estimated as follows. (a) Ni/Al 12 O 19 catalyst: At a superficial gas velocity of 6.7 m/s, the reactor length required to get the 99% of the equilibrium conversion of CO 2 is 0.18 m and the assumed pressure Results and discussion 121 drop of 1 bar is also just reached (Fig. 5-59). As this length is rather small, a length of 0.5 m can be assumed to be appropriate (to be on the safe side with regard to bypass effect etc.), even if the pressure drop then increases to 2.5 bar. The cross-sectional area of the RWGS reactor is 1.4 m 2 (ratio of volume rate of 9.4 m 3 /s and gas velocity of 6.7 m/s), the reactor diameter 1.3 m, and the volume of the catalyst is 0.7 m 3 . Fig. 5-59: Reactor length versus superficial gas velocity: dashed-dotted line: length, where a pressure drop of 1 bar is just reached; lines: length required to reach 99% of equilibrium conversion (X CO2 = 81%). (b) Al 2 O 3 catalyst: Now a lower superficial gas velocity of 1.7 m/s has to be adjusted to reach 99% of the equilibrium conversion of CO 2 and also a pressure drop of 1 bar (Fig. 5-59). The corresponding length of the fixed bed is about 2.8 m, and the crosssectional area 5.5 m 2 (2.6 m diameter). Increasing the reactor length by 1 m (total length of 3.8 m) to be on the safe side then corresponds to a volume of the Al 2 O 3 catalyst of about 21 m 3 , which is by a factor of 30 higher compared to the Ni-catalyst. These estimations clearly show that for both catalysts, the required volume is rather small to deliver the syngas for a huge technical FT reactor. The question, whether the Nior the Alcatalyst should be used is then a question of the costs of the catalyst and even more important of the long term stability of the catalyst, which should be determined by further investigations. 234 4 42 2     42 Reactor length [m] Superficial gas velocity, us[m/s] 5pb= 1 bar Al2O3 5pb= 1 bar Adiabatic reactor 5pb= 1 bar Ni/Al12O19 Zusammenfassung und Ausblick 128 In der vorliegenden Arbeit wurde der dritte Schritt, die RWGS-Reaktion, untersucht. Dabei wurde vor allem ein aluminiumoxidgeträgerter Nickelkatalysator (Ni/Al 12 O 19 ) verwendet. Da die RWGS-Reaktion endotherm verläuft, wurde die Reaktion bei hohen Temperaturen durchgeführt, um CO (und H 2 O) als Hauptprodukte zu erhalten. Übersicht zu den experimentellen Untersuchungen: Die experimentellen Untersuchungen aller Teilschritte (Methanisierung, Wasser-Gas-Shift bzw. RWGS) wurden in einem Festbettreaktor unter Verwendung des oben genannten Nickelkatalysators bei Normaldruck durchgeführt. Neben der Temperatur wurden auch die Konzentrationen der Reaktanten, die Verweilzeit und die Katalysatorpartikelgröße variiert. Für jede Reaktion erfolgte eine Bestimmung sowohl der intrinsischen als auch der effektiven Reaktionskinetik. Neben dem Nickelkatalysator wurden auch andere Katalysatoren wie z.B. Al 2 O 3 getestet. Die experimentellen Ergebnisse lassen sich folgendermaßen zusammenfassen. CO 2 -Hydrierung: Zunächst wurden verschiedene kommerzielle Katalysatoren untersucht (siehe Tab. 5-4), um ein grundsätzliches Verständnis für die Aktivität der Katalysatoren für die CO 2 -Hydrierung zu erhalten. Hierbei zeigte der Ni/Al 12 O 19 -Katalysator die höchste Aktivität. Daher wurden die weitergehenden kinetischen Untersuchungen zur CO 2 - Hydrierung mit diesem Katalysator durchgeführt. Erwartungsgemäß erhöhen sich der CO 2 -Umsatz und die CO-Ausbeute mit steigender Temperatur (300 – 950 °C). Bis zu einer Temperatur von 500 °C steigt zunächst auch die Methanausbeute an, um dann aufgrund der thermodynamischen Limitierung wieder abzufallen. Bei hohen Temperaturen wurden Umsätze erzielt, die sehr nahe am Gleichgewicht liegen. So konnte beispielsweise bei 950 °C ein Umsatz von 94 % des Gleichgewichtsumsatzes erhalten werden (leerrohrbezogene Verweilzeit: 18 ms). Die Partikelgröße wurde mit weniger als 0,5 mm so gewählt, dass der äußere Stofftransport keinen Einfluss hat. Der Nickelkatalysator zeigte bei 900 °C zudem eine sehr gute Langzeitstabilität. Die Berechnung der intrinsischen kinetischen Parameter ergab eine Aktivierungsenergie von 65 kJ/mol und einen Häufigkeitsfaktor von 5,55 · 10 2 m 1,5 mol 0,5 kg -1 s -1 . Die Reaktionsrate der CO 2 - Hydrierung kann durch folgenden Ansatz beschrieben werden: D E  | } ~ F  E  | } ~   |} ~   @    ~   r 1 2 t ¿ 04  Zusammenfassung und Ausblick 129 Die Berechnung des Stofftransporteinflusses auf die effektive Reaktionsrate der CO 2 -Hydrierung zeigt, dass die experimentellen und berechneten Ergebnisse gut übereinstimmen. Bei dem verwendeten Schalenkatalysator (Durchmesser 3 mm, Schichtdicke 0,5 mm) beeinflusst sowohl die Porendiffusion als auch der externe Stofftransport die effektive Reaktionsrate. Bei den experimentell eingestellten geringen Gasgeschwindigkeiten wird die Katalysatorleistung oberhalb von 900 °C zunehmend durch den externen Massentransfer bestimmt. CO-Hydrierung (Methanisierung): Die Methanisierung ist exotherm und somit bei niedrigen Temperaturen thermodynamisch begünstigt. Daher steigt der CO-Umsatz mit wachsender Temperatur an, um ab ca. 450 °C schließlich zu fallen. Zwischen 480 und 710 °C deaktiviert der Katalysator aufgrund der hier einsetzenden Koksbildung. Die Reaktionsordnungen wurden bei 340 °C und Atmosphärendruck bestimmt. Für CO ergibt sich ein Wert von -0,3 und für Wasserstoff von 0,7. Die Aktivierungsenergie beträgt 102 kJ/mol und der Häufigkeitsfaktor hat einen Wert von 2,35 · 10 5 m 1,2 mol 0,6 kg -1 s -1 . Die Berechnungen zum Stofftransport lassen für den Schalenkatalysator einen Einfluss der Porendiffusion oberhalb von 420 °C erwarten. Die berechneten Werte der effektiven Reaktionsrate weichen allerdings bei hohen Temperaturen von den experimentellen Werten ab. Dies ist auf die dann auftretende Deaktivierung des Katalysators zurückzuführen. Wasser-Gas-Shift Reaktion (WGS): Die WGS-Reaktion wurde ebenfalls untersucht, da es sich hier um die Rückreaktion der CO 2 -Hydrierung (RWGS) handelt. Die Hauptprodukte der WGS-Reaktion sind daher Wasserstoff und Kohlendioxid. Diese Reaktion ist exotherm und wird folglich bei niedrigen Temperaturen thermodynamisch begünstigt. Bis zu einer Temperatur von 450 °C erhöhen sich sowohl der CO-Umsatz als auch die CO 2 - Ausbeute. Gleichzeitig konnte in dem untersuchten Temperaturfenster weder eine Methanbildung noch eine Deaktivierung des Katalysators beobachtet werden. Im Temperaturbereich von 500 bis 700 °C zeigt der Nickelkatalysator eine Deaktivierung, die durch eine Verkokung des Katalysators hervorgerufen wird. Die Zugabe beider Reaktanten (CO, H 2 O) übt einen positiven Einfluss auf die WGSReaktion aus, während (wie zu erwarten) die Beimischung der Produkte (CO 2 , H 2 ) einen negativen Einfluss auf die Reaktionsrate aufweist. Die Bestimmung der Reaktionsordnung für die einzelnen Komponenten ergab: 0,8 (CO), 0,4 (H 2 O), -0,1 (CO 2 ) und -0,15 (H 2 ). Die Bildung von Methan fand im untersuchten Parameterbereich nur untergeordnet statt. Zusammenfassung und Ausblick 130 Die Aktivierungsenergie der WGS-Reaktion (Ni-Katalysator) beträgt 96 kJ/mol. Der Wert des Häufigkeitsfaktors liegt bei 3,46 · 10 5 m 3,6 mol -0,2 kg -1 s -1 . Stofftransportberechnungen für den Schalenkatalysator zeigen oberhalb von 350 °C einen Einfluss der Porendiffusion. Bedingt durch eine Katalysatordeaktivierung nimmt die effektive Reaktionsrate im Temperaturbereich zwischen 500 und 700 °C ab. Oberhalb von 900 °C wird die effektive Reaktionsgeschwindigkeit mehr und mehr durch den externen Massentransfer bestimmt. Zur Modellierung eines technischen Reaktors zur Herstellung von Synthesegas durch die CO 2 -Hydrierung wurde für beide Katalysatoren (Ni/Al 12 O 19 ; Al 2 O 3 ) das Modell eines eindimensionalen Festbettreaktors verwendet, d.h. die axialen Temperaturund Umsatz-Profile wurden berechnet. Dabei wurde der adiabate und auch der isotherme Fall berücksichtigt. Die effektive Reaktionsrate der RWGS-Reaktion am Ni/Al 12 O 19 -Katalysator ist im Vergleich zum Al 2 O 3 -Katalysator um den Faktor 40 höher. Für einen gegebenen Umsatz kann die Reaktorlänge für den Ni-Katalysator entsprechend wesentlich kürzer gewählt werden. Ausblick Ausgehend von den in dieser Arbeit erhaltenen Ergebnissen sollten in weiterführenden Untersuchungen sowohl für den Ni-Katalysator als auch für den Al 2 O 3 -Katalysator ihre Langzeitstabilitäten untersucht werden. Hierzu sind Versuchszeiten von einigen Wochen anzustreben, um aussagekräftige Daten zu generieren. Weiterhin sollten Temperaturen von mehr als 1000 °C eingestellt werden, da dann der erreichbare CO 2 -Umsatz bei der RWGS allein schon aus thermodynamischen Gründen noch höher ist (> 90%). Hierzu muss die bestehende Versuchsanlage mit einem Reaktor, der diesen Temperaturen standhält, sowie einem entsprechenden Heizsystem ausgestattet werden. Appendix A 131 Appendix A A.1 Utilization of CO 2 in catalytic conversion processes The detailed description for the synthesis of value added products using CO 2 is given in this section. A.1.1 Salicylic acid formation The direct synthesis of salicylic acid via the coupling of CO 2 with phenol is a “green process”, useful for the chemical fixation of CO 2 . Salicylic acid is widely used in the industry as a raw material and intermediate for the production of pharmaceutical and fine chemicals.  x j ¡ ij  <   i P  À   x j r 2 ij 4 iij 1 2g ¿ 04  Salicylic acid can be obtained by carboxylation of phenol with a carbon dioxide in presence of Lewis acid catalysts. AlBr 3 as a catalyst showed the best activity and selectivity towards salicylic acid among the various Lewis acids examined [108]. At 80 °C and 8 MPa, the yield of salicylic acid reached to 60% with the selectivity of 100% in 1 h. Several types of base metal oxides such as alumina, zirconia, ceria, and alkali metal and alkaline earth metal salts have been used for the synthesis of salicylic acid [109, 110]. Potassium carbonate as a catalyst under optimum condition yields 68% of salicylic acid with 99% selectivity. The systematic investigation for the direct reaction of supercritical CO 2 and phenol over the catalysts like ZrO 2 , TiO 2 , and KF/NaY+Mn 2+ , ZrO 2 was investigated [1]. A.1.2 Synthesis of urea and urea derivatives The process of urea synthesis from carbon dioxide and ammonia was developed in 1922, and is called the Bosch-Meiser urea process. This non-catalytic process still exists and operates at relatively high pressure and high temperature. = Á j u  <   i P  À  j P ÁiiÁ j r 1 2g ¿ =4  j P ÁiiÁ j r  À  Á j P i Áj P <  j P i 1 2g ¿ 34  Appendix A 132 The production of urea involves formation of ammonium carbamate (NH 2 COONH 4 ) using ammonia and CO 2 . Ammonium carbamate subsequently gets dehydrated into urea and water by the application of heat. The first reaction is fast, exothermic, and basically goes to completion, while the second one is relatively slow, endothermic, and does not go to completion [111]. Both of these reactions are reversible and therefore ammonia and carbon dioxide exit the reactor along with ammonium carbamate and urea. Thus, the conversion based on CO 2 is usually in the order 50 to 80% [111]. The components of this mixture in the product stream are then separated by stripping off gaseous ammonia followed by carbon dioxide. The worldwide urea production in 2008 was 146 million metric tons which corresponds to 107 million metric tons of CO 2 consumption per annum [112]. Hence in the chemical industry, only this process has a certain potential of CO 2 reduction. Urea is mainly used for fertilizers but urea derivatives are used in the various other areas such as agrochemicals, pharmaceuticals, antioxidants in gasoline, and corrosion inhibitors [113]. Commercially urea derivatives are synthesised by using toxic phosgene as well as carbon monoxide. In a new approach, urea derivatives are produced by using amines and CO 2 as a carbonyl agent via catalytic and non-catalytic processes. The carbonylation of amines using CO 2 is simple, safe, and a clean process which avoids the use of poisonous compounds like phosgene and CO [114, 115]. The reaction of CO 2 with primary aliphatic amines also leads to urea derivative [114]. The process using the ionic liquid [Bmim]Cl with CsOH as catalyst yields about 98% of an urea derivative [115]. A.1.3 Styrene synthesis Commercially styrene is produced by ethylbenzene dehydrogenation using potassium promoted iron oxide as a catalyst with a large excess of superheated steam at 600 to 650 °C. The steam plays an important role in ethylbenzene dehydrogenation as it shifts the equilibrium towards higher conversions, decreases the amount of coke formation and supply the heat for the reaction. A new route uses carbon dioxide for the dehydrogenation: Appendix A 133 Conventional route: 1 2g¿p4 New CO 2 -based route: 1 2g¿`4 The catalytic conversion of ethylbenzene using CO 2 have been discussed recently by Park et al. [116]. The direct utilization of CO 2 as an oxidant in the dehydrogenation of ethylbenzene to styrene offers several advantages. The main advantages shown in different studies [117-119] are the acceleration of the reaction rate, the drop in temperature, the enhancement of the conversion of ethylbenzene and of the selectivity. Moreover, the energy consumption for the ethylbenzene dehydrogenation using carbon dioxide is much lower than the currently operating process using steam [116]. The dehydrogenation of ethylbenzene to styrene in the presence of CO 2 was also studied over MnO 2 -ZrO 2 [117], SnO 2 –ZrO 2 mixed oxide nano-composite [118], cromia based catalyst [119], Co, Mo and CoMo catalysts supported on natural and aluminium-pillared clays [120], and a CeO 2 –ZrO 2 mixed oxide supported SBA-15 catalyst [121]. In these catalytic systems, acidic and basic sites cooperatively activate CO 2 as well as ethylbenzene, which lead to a higher conversion of ethylbenzene. An attempt to use CO 2 as a diluent and oxidant over an activated carbon-supported iron catalyst have made by Sugino et al. [122]. The addition of 20 to 30 mol% lithium nitrates to the iron catalyst resulted in a significant increase in the catalytic activity. A.1.4 Dimethyl carbonate (DMC) synthesis The conventional DMC synthesis use toxic phosgene. The use of carbon dioxide in the synthesis of DMC presents an environmentally friendly and attractive approach since it replaces phosgene and chlorine [123, 124]. Appendix A 134 Conventional route:  i + P  <  = j u ij  v  j u iii j u  <  = j+  2g ¿ Y4  New CO 2 -based route: i P  <  = j u ij  v  j u iii j u  <  j P i  2g ¿ t4  DMC has been synthesized using CO 2 and methanol over solid oxide catalysts like ZrO 2 [125, 126]. The dissociative adsorption of CO 2 occurs faster than the adsorption of methanol on ZrO 2 and species formed from methanol are bound more strongly. DMC also has been synthesized using base catalysts like K 2 CO 3, Na 2 CO 3, Cs 2 CO 3 and CH 3 I as a promoter [127]. The use of an ionic liquid along with CH 3 I as a promoter over KOH catalyst shows an increase in the yield of DMC [128]. The utilization of CO 2 for synthesis of DMC is very beneficial because DMC itself is a unique molecule and can be used in new environmentally friendly reactions to replace some of the environmentally harmful processes like transesterification of DMC with phenol that produces methyl phenyl carbonate [129]. A.1.5 Methanol synthesis Methanol is used as an important chemical feedstock for the synthesis of many compounds such as acetic acid, formaldehyde, methyl tert-butyl ether (MTBE), methyl methacrylate (MMA) and chloromethane. The first industrial plant for methanol production using synthesis gas was constructed by the BASF in 1923. This process was operating under high pressure of about 20 MPa and at 300 °C using a zinc oxide/chromium oxide catalyst [5]. Due to the environmental impact, methanol synthesis through the carbon dioxide hydrogenation has attracted worldwide research interest. The potential use of CO 2 to replace CO in the methanol synthesis is also an effective way of CO 2 utilization [130]. The methanol synthesis from CO 2 has two possible mechanisms. One is the direct conversion of carbon dioxide to methanol (Eq. (A-8)) and the second one is the conversion of CO 2 to CO and H 2 O via the RWGS reaction with subsequent conversion of CO and H 2 to methanol (Eq. (A-9)). Commercially methanol has been produced from coal and natural gas containing small amount of CO 2 along with CO and H 2 as a feedstock [131]. i P  <  3 j P  v  j u ij  <  j P i       1 j l Pmn  o  F  5 p\ [ `K  q O -+  2g ¿ K4  Appendix A 135 i  <  = j P  v  j u ij  1 j l Pmn  o  F  5 \Z [ tt  q O -+  2g ¿ \4  The methanol formation is an exothermic reaction, so thermodynamically favored at low temperatures and high pressures. During the reaction hydrogen can be consumed by CO 2 in reverse water gas shift reaction and reduce the methanol formation. A recent study show that a mixture of a proper proportion of CO 2 and CO for methanol synthesis is not only increases the methanol yield but also decrease the apparent activation energy of the reaction [132]. The presence of CO 2 could maintain the active copper sites in the oxidation state or prevent an over-reduction of the ZnO component of Cu/Zn catalysts during methanol synthesis [132]. Table A-1 shows the data reported for conversion of CO 2 and the selectivity to methanol using CO 2 over Cu/Zn/ZrO 2 containing catalyst. Table A-1: Methanol synthesis by CO 2 hydrogenation using Cu/Zn/ZrO 2 containing catalyst. Catalyst Preparation method T, °C X CO2 , % S MeOH , % Ref. Cu/Zn/ZrO 2 coprecipitation 250 19.4 29.3 [133] Cu/Zn/ZrO 2 coprecipitation 220 21 68 [134] Cu/Zn/ZrO 2 coprecipitation 200 5.9 47.5 [135] Cu/Zn/ZrO 2 urea nitrate combustion 240 17 56.2 [136] Cu/Zn/ZrO 2 glycine nitrate combustion 220 12 71.1 [137] Cu/Zn/Al/ZrO 2 coprecipitation 240 18.7 47.2 [138] Cu/Zn/Ga/ZrO 2 coprecipitation 250 - 75 [139] The presence of ZrO 2 enhances Cu dispersion which leads to an improved catalytic activity and stability towards methanol formation [133, 136]. The addition of an optimum amount of metal oxides like Ga 2 O 3 , Al 2 O 3 , ZrO 2 , and Cr 2 O 3 as a modifiers increases the methanol synthesis activity and the stability of Cu/ZnO-based ternary catalysts [135, 138-140]. A novel approach was employed using low temperature of about 170 °C and 5 MPa pressure in a semi-bath autoclave reactor with 2-butanol as a solvent [141]. CO 2 conversion and Appendix A 136 selectivity towards methanol obtained over a Cu catalyst was about 26% and 73%, respectively. Although much attention has been given to the methanol synthesis from CO/H 2 , CO 2 /H 2 and CO/CO 2 /H 2 using fixed bed technology, some important subjects of investigation are still under discussion: (1) whether the active site is Cu 0 or Cu n+ , (2) whether the carbon source for methanol synthesis is CO 2 or CO, (3) whether there is special interaction between the Cu metal and oxide support, and (4) whether the water formed during methanol synthesis has any influence on the catalytic activity. A.1.6 Formic acid synthesis The synthesis of formic acid is another alternative to convert carbon dioxide into valuable products. Formic acid is widely used as a pickling agent, as a reducing agent, as an antibacterial agent, a mordant in the dyeing industry, disinfectant and preservative agent in sanitary stations, and as a neutralizer in the tanning industry [142]. It has been also used as a raw material in the chemical industry for the production of formate ester, which is important feedstock for the synthesis of various organic derivatives like aldehydes, ketones, amides, and carboxylic acids [142]. Formic acid has been produced directly by the hydration of carbon monoxide, and by the hydrolysis of methyl formate. The synthesis of formic acid by direct CO 2 hydrogenation was first reported by Farlow and Adkins in 1935 using Raney nickel as a catalyst under 20 to 40 MPa and 80 to 150 °C in presence of amines [143]. At a temperature of 250 °C, the sheet brass used in fabricating liners for steel reaction vessel acts as an active catalyst for the hydrogenation reaction: i P  <  j P  À jiij 1 2g ¿ 0Z4  Most of the catalytic studies of formic acid synthesis are based the metal complexes of the second and third row transition metals, where the metal like ruthenium, palladium, and rhodium has been used with the combination of halides or hydrides as anionic ligands and phosphines as neutral ligands [142]. The catalytic hydrogenation of CO 2 to formic acid over transition metal catalyzed complexes RuH 2 (PPh 3 ) 4 and Pd(Ph 2 PCH 2 CH 2 PPh 2 ) 2 , has been reported [144]. Similarly, formic acid has been produced by carbonylation of CO and H 2 O using a metal complex, i.e. Appendix A 137 [Ru II (EDTA-H)CO] - in the WGS reaction [145]. The highly active metal complex RuCl(O 2 CMe)(PMe 3 ) 4 as a catalyst with acidic alcohols was used to enhance the rate of the hydrogenation reaction. The organic bases with intermediate basicity were added to the reaction mixture to extract the formic acid [146]. Although much attention has been paid to synthesize formic acid using a homogeneous catalyst, the problem of separation of formic acid from the catalyst and the base still remains. The hydrogenation of CO 2 to formic acid over alumina supported ruthenium hydroxide as a heterogeneous catalyst was carried out using triethylamine and ethanol as a solvent at 80 °C and 13.5 MPa [142]. Also the Ruthenium(II) complex catalyzed hydrogenation of CO 2 to formic acid was theoretically investigated using cis-RuH 2 (PH 3 ) 4 as a model catalyst [147, 148]. The hydrogenation over this catalyst took place through the insertion of CO 2 into the Ru-H bond followed by the H-OCOH reductive elimination, where CO 2 insertion was the rate determining step. A.1.7 CO 2 reforming of methane The reforming of CO 2 with methane has also attracted a continuous research interest. This process was carried out not only for the utilization of two undesirable greenhouse gases (CO 2 , CH 4 ) but also for the production of fuels by Fischer Tropsch and methanol synthesis. The dry reforming of methane with CO 2 is not yet feasible, and gives incomplete conversion of CO 2 due to thermodynamic constraints [149]. Dry reforming: i P  <  j r  v  = i  <  = j P       1 j l Pmn  o  F  =pt  q O -+  2g ¿ 004  Steam reforming: j r  <  j P i  v  i  <  3 j P       1 j l Pmn  o  F  =ZY  q O -+  2g ¿ 0=4  The dry reforming of CO 2 with methane was first studied by Fischer and Tropsch in 1928. After that many researchers investigated this reaction using various catalysts for the production of syngas. Syngas can be produced with CO/H 2 ratio of about a unity depending upon the reaction conditions. This is an endothermic reaction, generally carried out in the temperature range of 300 to 830 °C and at atmospheric pressure [150]. The group VIII metals (e.g. Ni, Co, Pt, Pd, Ir, Ru, Rh etc.) are more or less catalytically active for the carbon