- Email: [email protected]
Kinetics studies of nano-structured cobalt–manganese oxide catalysts in Fischer–Tropsch synthesis
Journal of Industrial and Engineering Chemistry 19 (2013) 1177–1183
Contents lists available at SciVerse ScienceDirect
Journal of Industrial and Engineering Chemistry journal homepage: www.elsevier.com/locate/jiec
Kinetics studies of nano-structured cobalt–manganese oxide catalysts in Fischer–Tropsch synthesis Mohsen Mansouri a, Hossein Atashi a,*, Farshad Farshchi Tabrizi a, Ali Akbar Mirzaei b, Ghobad mansouri c a
Department of Chemical Engineering, University of Sistan & Baluchestan, Zahedan 98164-161, Iran Department of Chemistry, University of Sistan & Baluchestan, Zahedan 98135-674, Iran c Department of Chemistry, University of Payame Noor, Ilam 69319-36173, Iran b
A R T I C L E I N F O
Article history: Received 24 September 2012 Accepted 11 December 2012 Available online 20 December 2012 Keywords: [Co(NH3)4CO3]MnO4 precursor Cobalt–manganese catalyst Fischer–Tropsch synthesis Kinetic model
A B S T R A C T
The nano-structured cobalt/manganese oxide catalyst was prepared by thermal decomposition of [Co(NH3)4CO3]MnO4 precursor, and was tested for the Fischer–Tropsch reaction (hydrocarbon forming) in a fixed-bed micro-reactor. Experimental conditions were varied as follow: reaction pressure 1–10 bar, H2/CO feed ratio of 1–2 and space velocity of 3600 h1 at the temperature range of 463.15–523.15 K. On the basis of carbide and/or enolic mechanisms and Langmuir–Hinshelwood–Hougen–Watson (LHHW) type rate equations, 30 kinetic expressions for CO consumption were tested and interaction between adsorption HCO and dissociated adsorption hydrogen as the controlling step gave the most plausible kinetic model. The kinetic parameters were estimated with non-linear regression method and the activation energy was 80.63 kJ/mol for optimal kinetic model. Kinetic results indicated that in Fischer– Tropsch synthesis (FTS) rate expression, the rate constant (k) has been increased by decreasing the catalyst particle size. The catalyst characterization was carried out using different methods including powder X-ray diffraction (XRD), scanning electron microscopy (SEM) and Brunauer–Emmett–Teller (BET) surface area measurements. ß 2012 The Korean Society of Industrial and Engineering Chemistry. Published by Elsevier B.V. All rights reserved.
1. Introduction Nowadays, there is a renewed interest in Fischer–Tropsch synthesis (FTS), especially for the selective production of petrochemical feedstock such as ethylene, propylene and butylene (C2–C4 olefins) directly from synthesis gas [1,2]. Due to the thermodynamic and kinetic limitations of the reaction, it is believed that bimetallic catalysts system containing one or two metal active agents increase the value of the light olefins whereas catalyst composition exerts the greatest influence on the molecular weight distribution for the FTS products [3,4]. It is shown that the addition of Mn to Fe or Co catalysts leads to a significant increase in high olefin formation and a decrease in methane activity [5,6]. Co–Mn catalysts have been investigated intensively for its higher selectivity to lower molecular weight olefins [7,8]. The kinetic description of the FT reaction is extremely important for the industrial practice, being a prerequisite for the industrial FT process design, scale-up, optimization, and simulation. Intrinsic reaction rate equations for the FT equation
* Corresponding author. E-mail address: [email protected] (H. Atashi).
(hydrocarbon forming only) based on Langmuir–Hinshelwood– Hougen–Watson (LHHW) adsorption theory have been developed and used successfully for iron- and cobalt-based catalysts [9–21]. The mechanisms proposed include the carbide, enolic, and direct insertion theories [22]. An investigation was undertaken to elucidate the reaction kinetics of the Co–Mn oxide catalysts used by Keyser et al. [9] and to compare its performance with other cobalt- and iron-based catalysts. It was found that a reaction rate equation for FT reaction based on the enolic theory performed slightly better than a reaction rate equation based on the carbide theory. Chen and Adesina [23] have studied improved alkene selectivity over a silica supported cobalt–molybdenum catalyst, and their data also suggest the carbide mechanism where the rate determining step is the dissociative adsorption of hydrogen. Sari et al. [10] studied the rate of FTS over alumina-supported Co–Ru in a slurry phase reactor. They considered five parametric kinetic expressions based on their previous studies: one empirical power law model and four variations of the LHHW approach regarding strong inhibition due to CO adsorption and the kinetic data, fitted by a simple LHHW. Although numerous kinetic expressions were developed for Co-based catalyst, one should consider the kinetics of bimetallic
1226-086X/$ – see front matter ß 2012 The Korean Society of Industrial and Engineering Chemistry. Published by Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.jiec.2012.12.015
M. Mansouri et al. / Journal of Industrial and Engineering Chemistry 19 (2013) 1177–1183
1178
2.2. Catalyst characterization Nomenclature r FT a b k K PCO P H2 0 FCO R T W X r CO s
uS uHOC uH E
DH N
(mol g1 cat
1
rate of reaction min ) adsorption parameter (see Table 4) adsorption parameter (see Table 4) reaction rate constant (see Table 4) adsorption constant (see Table 4) CO pressure (bar) hydrogen pressure (bar) molar flow rate of CO at the inlet (mol min1) gas constant (8.314 J mol1 K1) temperature (K) catalyst mass (g) conversion the consumption rate of CO unoccupied active sites the concentration of free sites surface occupied with HOC surface occupied with H activation energy (kJ/mol) heat of adsorption (kJ/mol) number experimental data points
Superscripts experimental value exp conditions in feed in predicted value cal
Co-based FT catalyst. Thus, the main objective of the present work is to investigate the kinetic and mechanism of the CO hydrogenation on nano-structured Co–Mn catalyst, which was prepared by thermal decomposition of [Co(NH3)4CO3]MnO4 precursor. We establish kinetic models for hydrocarbon formation on the basis of the correlation between experimental data and supposed reaction mechanism sets. The suitable model was obtained and the kinetic parameters were appointed. The catalyst characterization was carried out using different methods including XRD, SEM and BET. 2. Experimental
The catalyst was analyzed by X-ray diffraction and all peaks were consistent with the peaks standard Co–Mn oxide. No peaks from any other phases of Co–Mn oxide were observed. The XRD pattern when compared with JCPDS reveals the structure as spinel oxide and a mixed oxide spinel structure (CoMn) (CoMn)2O4 was identified which is similar to that reported by van der Reit [13]. The morphology of catalysts and their precursors was observed by means of an S-360 scanning electron microscopy (USA). BET surface areas, pore volumes and average pore sizes of the catalyst precursor and calcined samples were measured by N2 physisorption using a Quantachrome Nova 2000 automated system (USA). Each catalyst sample was degassed under nitrogen atmosphere at 300 8C for 3 h. In order to obtain the BET surface areas, pore volumes and average pore sizes, different samples were evacuated at 196 8C for 66 min. 2.3. Catalyst testing The kinetic experiments were carried out in a tubular fixed bed micro-reactor. A schematic representation of the reactor is shown in Fig. 1. The reactor consisted of a single stainless steel tube with an inner diameter of 20 mm surrounded by an alumina jacket to achieve a uniform wall temperature along the length of the reactor. A preheating zone a head of catalyst packing was filled with inert quartz glass beads. External heating was provided by an electrical element wrapped around the alumina jacket and placed in through the fire brick part. The required amount of the catalyst is diluted by asbestos and placed among the inert quartz glass beads. The temperatures of all of different zones include: preheating zone, catalyst bed and underneath zone of the reactor is checked by three separate thermocouples placed in different parts of the reactor. The thermocouple which controls bed temperature is placed exactly below the catalyst bed. The inlet feed gas is arrived from top of the reactor and the outlet products are avoided from underside of the reactor. All gas lines to the reactor bed were made from 1/400 stainless steel tubing. Three mass flow controllers (Brooks, Model, 5850E) were used to adjust automatically flow rate of the inlet gases comprising CO, H2 and N2 (purity of 99.99%). Mixture of CO, H2 and N2 were subsequently introduced into the reactor, which was placed inside a tubular furnace (Atbin, Model ATU 150-15). Prior to the reaction the catalysts were reduced in situ using H2 (30 ml/min) and N2 (30 ml/min) mixture gas at 350 8C for 16 h. In
2.1. Catalyst preparation and characterization The catalyst used in the present study was prepared using a thermal decomposition precursor as follows. An aqueous solution (60 ml) of ammonium carbonate (20 g) was added to a concentrated aqueous NH3 (60 ml), and the mixture was stirred at room temperature for 30 min. To this solution was then added an aqueous solution (30 ml) of Co(NH3)26H2O (15 g) and the resulting solution was stirred at room temperature for 1 h. Then, 8 ml of 30% hydrogen peroxide was dropwise added into the solution, stirring continuously. The final solution was filtered and left for slow evaporation in air until water-soluble [Co(NH3)4CO3]NO3 crystals were obtained. To a solution containing [Co(NH3)4CO3]NO3 (1 mmol) in 100 ml water was added KMnO4 (1 mmol), and the mixture was stirred for several minutes. The precipitates ([Co(NH3)4CO3]MnO4 precursor) were recovered by centrifugation and washed by distilled water and then dried in ambient air at 30 8C overnight. In order to obtain the final nanocatalyst, the precursor was calcined at 400 8C in static air in the electric furnace for 4 h. The gray powder, nano-structured Co–Mn oxide catalyst was formed and kept in desiccator.
Fig. 1. Different parts of micro-fixed-bed reactor.
M. Mansouri et al. / Journal of Industrial and Engineering Chemistry 19 (2013) 1177–1183
each test, 1.0 g catalyst was loaded and the reactor operated about 12 h to ensure attaining the steady state operating conditions. The catalysts were extremely fine particles so intraparticle diffusion could be neglected. The gas hourly space velocity (GHSV) increased to the value in which the CO conversion was almost the same for a variety of catalyst weight which indicates that external diffusion can be neglected above this GHSV. Hence, the kinetic experiments were conducted free from internal and external mass transfer limitations. Experiments were carried out with mixtures of H2, CO and N2 in a temperature range of 463.15–523.15 K, H2/CO feed ratio of 1/1–2/1, pressure range of 1–10 bar and GHSV = 3600 h1. The catalytic test data are represented in Table 1. The stability of the catalyst was investigated by repetition of central points of the designs in the middle and the end of the experiments. We have a differential flow reactor when we choose to consider the rate to be constant at all points within the reactor. Since rates are concentration-dependent this assumption is usually reasonable only for small conversions or for shallow small reactors. For conversions up to 15%, the curve conversion versus inverse rate is nearly constant; hence the assumption of differential reactor can be used. For each run in a differential reactor, the plug flow performance equation becomes as follows: W ¼ 0 FCO
Z
xout xin
dx xout xin xout 0 ¼ ¼ r CO r CO r CO
(1)
According to the above equation, the average rate for each run is derived as follow: r CO ¼
1179
several agglomerates of irregularly spherical shaped grains and shows that this material has a less dense and homogeneous morphology. After the calcination at 400 8C for 4 h, the morphological features are different with the precursor sample and show that the agglomerate size is greatly reduced compared to the precursor (Fig. 2b). It may be due to that the calcined catalyst surface is covered with small crystallite of cobalt and manganese oxide. After FTS chemical reaction the catalyst texture and its morphology changed (Fig. 2c). However, the size of these grains grew larger by agglomeration in the tested catalyst, which may be due to the sintering after reactions. 3.1.2. BET measurements BET surface area, pore volume and pore diameter of the precursor, nano and conventional (prepared using co-precipitation method by Keyser et al. [9]) catalysts are summarized in Table 2. As shown from this Table, the BET surface area and pore volume of precursor are low, about 11.32 m2/g and 0.12 cm3/g, respectively. It can be seen that after calcinations, both the BET surface area and pore volume increased to 74.28 m2/g and 0.21 cm3/g, respectively; this is in agreement with the SEM results which showed that the agglomerate size of calcined catalyst was less compared with its precursor and therefore leaded to an increase in the BET specific surface area of the calcined sample. Also, it can be found that the BET surface area of the nano catalyst is much higher than the conventional catalyst used by Keyser et al. [9]. 3.2. Kinetic modeling and reaction rate equations
0 FCO xout W
(2)
3. Result and discussion 3.1. Catalyst characterization 3.1.1. Morphological properties The scanning electron microscopy (SEM) observations have shown differences in morphology of precursor and calcined catalysts (before and after the reaction) and indicate that these materials are made of nano-metric particles. The electron micrograph obtained from catalyst precursor (Fig. 2a) depicts
Although great efforts have been devoted to elucidate the FTS reaction mechanism, there are still many controversies on this point, and many theories have been proposed in the literature [14– 16]. On the basis of the nature of CO adsorption and of the nature of chain initiator intermediates, popular mechanistic proposals include (1) the carbide mechanism [17,22], wherein CO adsorbs dissociatively and the carbide (Cs) is the chain initiator intermediate. In this mechanism simultaneous dissociative adsorption of CO and H2 is followed by the hydrogenation of adsorbed carbon by adsorbed hydrogen in a stepwise manner to give methane and higher hydrocarbons. (2) The enolic mechanism [18,22], involving molecular adsorption of CO which reacts with adsorbed hydrogen
Table 1 Summary of experimental conditions and results for kinetic tests at PTot. = 1–10 bar, T = 463.15–523.15 K and GHSV = 3600 h1 in a fixed bed reactor (FBR). No.
T (K)
H2/CO
P CO (bar)
P H2 (bar)
0 FCO 103 (mol min1)
XCO (%)
1 r CO 102 (mol g1 ) cat min
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22
463.15 463.15 473.15 488.15 498.15 498.15 498.15 463.15 463.15 488.15 488.15 493.15 493.15 498.15 498.15 463.15 463.15 468.15 488.15 498.15 503.15 523.15
1 1 1 1 1 1 1 1.5 1.5 1.5 1.5 1.5 1.5 1.5 1.5 2 2 2 2 2 2 2
0.9 2.1 1.5 0.9 0.9 1.5 2.1 0.72 1.2 1.2 1.68 0.72 1.2 1.68 0.72 1.4 2 0.6 2 0.6 1.4 2
0.9 2.1 1.5 0.9 0.9 1.5 2.1 1.08 1.8 1.8 2.52 1.08 1.8 2.52 1.08 2.8 4 1.2 4 1.2 2.8 4
1.402 3.258 2.288 1.330 1.304 2.173 3.042 1.119 1.858 1.774 2.484 1.058 1.745 2.434 1.043 2.181 3.116 0.925 2.957 0.867 2.008 2.759
8.448 9.28 7.7 8.58 9.12 9.84 11.82 9.47 9.69 12.48 11.79 12.75 12.36 11.91 13.28 9.615 9.762 8.93 11.39 12.5 12.17 14.21
1.185 3.023 1.762 1.142 1.188 2.138 3.595 1.060 1.800 2.214 2.928 1.349 2.157 2.899 1.385 2.097 3.042 0.871 3.369 1.143 2.444 3.920
M. Mansouri et al. / Journal of Industrial and Engineering Chemistry 19 (2013) 1177–1183
1180
Fig. 2. SEM micrographs of: (a) precursor, (b) calcined catalyst before the test, and (c) calcined catalyst after the test.
and the formation of an oxygenated intermediate, the enol (HOCs). In this mechanism, dissociative adsorption of H2 and molecular adsorption of CO followed by the hydrogenation of adsorbed carbon monoxide by adsorbed hydrogen to form an oxygenated intermediate which reacts with another adsorbed hydrogen to form water and adsorbed carbon, and the reaction of the resulting carbon with adsorbed hydrogen as in the carbide mechanism. To obtain the rate equation, a reaction mechanism should be adopted. Based on the elementary reactions explained in the above, we defined ten different possible mechanisms. The elementary reactions, which have been settled on sites for different models are listed in Table 3. These mechanisms were offered on the basis of various monomer and carbon chain distribution pathway. These mechanisms were obtained by considering different possibility of adsorption of CO and H2 molecules (associative or dissociative mode) on the catalyst surface. Each of these adsorption modes leads to different mechanisms and hence different kinetic equations were obtained. In order to derive rate equations, the LHHW approach was utilized and for each model, the possible ratedetermining steps were identified, while all other steps were assumed to be at quasi-equilibrium. Then, all of the obtained models were fitted separately, against experimental data. Table 4 displays the final form of different rate expressions for the 10 kinetic models. The development of the kinetic equations will be illustrated for model FT-IV3 (the third step elementary reaction from the FT-IV model). The model codes refer to the set of elementary reactions and the elementary reaction is not at equilibrium (that is the rate-determining step, so in this case reaction (3)). The set of elementary reactions for model FT-IV3 is shown in Table 3. Firstly, dissociated hydrogen reacted with molecular CO yielding formyl (HOC) intermediates (step 2) and these species were hydrogenated in steps 3 and 4, forming methyl monomers. The reaction rate of the rate-determining step is: r FT-IV3 ¼ k3 uHOC uH
(3)
where uHOC is the surface fraction occupied with the formyl intermediate and uH is the surface fraction occupied with the dissociative adsorbed hydrogen. The fraction of vacant sites, uS, can be calculated from the following balance equation:
quasi-equilibrium: k1
H2 þ 2s !2Hs 2
(5)
2
k1 P H2 uS k1 uH ¼ 0
uH ¼ K10:5 PH0:52 uS K1 ¼
(6)
k1 k1
where K1 is the equilibrium constant of dissociated hydrogen adsorption step. k2
CO þ Hs !HOCs
(7)
k2 P CO uH k2 uHOC ¼ 0
uHOC ¼ K 2 PCO uH ¼ K10:5 K 2 PH0:52 PCO uS K2 ¼
(8)
k2 k2
Substituting Eqs. (6) and (8) into Eq. (4), the concentration of free active site can be expressed as:
uS ¼
1 0:5 þ K P 0:5 1 þ K10:5 PH 2 CO PH2 2
(9)
with substituting of Eq. (9) into Eqs. (6) and (8), the expressions of [uH] and [uHCO] becomes:
uH ¼
1þ
uHOC ¼
0:5 K10:5 PH 2 0:5 0:5 0:5 K1 PH2 þ K 2 PCO PH 2
(10)
0:5 K10:5 K 2 PH P CO 2
(11)
0:5 þ K P 0:5 1 þ K10:5 PH 2 CO PH2 2
By substituting of the surface fraction of HOC and H in Eq. (3), the final rate expression is obtained as follows:
r FT-IV3 ¼
k3 K 1 K 2 PCO P H2 0:5 þ K P 0:5 ð1 þ K10:5 PH 2 CO PH2 Þ 2
2
¼
kPCO P H2 2
0:5 þ bP 0:5 ð1 þ aPH CO PH2 Þ 2
(12)
uS þ uH þ uHOC ¼ 1
(4)
The surface fractions of HOC and H can be calculated from the site balance, the preceding reaction steps which are at
Table 4 summarizes the final form of the different rate expressions for the 30 possible kinetic models, whereas Table 5 shows the kinetic and adsorption parameters for the several kinetic
Table 2 Textural of the precursor, nano and conventional catalysts. Sample
BET surface area (m2/g)
Pore volume (cm3/g)
Average pore diameter (nm)
[Co(NH3)4CO3]MnO4 Nano catalyst Conventional catalyst [9]
11.32 74.28 4.3
0.12 0.21 NA
16.89 13.32 NA
M. Mansouri et al. / Journal of Industrial and Engineering Chemistry 19 (2013) 1177–1183
1181
Table 3 Elementary reactions sets for FT synthesis. Model
Number
Elementary reaction
Model
Number
Elementary reaction
FT-I
1 2 3 4 5
CO + s $ COs H2 + 2s $ 2Hs COs + Hs $ CHOs + s CHOs + Hs $ CHOHs + s CHOHs + s $ CHs + OHs
FT-VI
1 2 3
H2 + S $ H2s CO + H2s $ CHOHs CHOHs + H2s $ CH2s + H2Os
FT-VII
1 2 3 4 1 2 3 1 2 3 4
CO + s $ COs H2 + s $ H2s COs + H2s+ $ CH2Os + s CH2Os+H2s $ CH2s+H2O CO + s $ COs COs + H2 $ CHOHs CHOHs + H2 $ CH2s + H2O H2 + 2s $ 2Hs CO + Hs $ COHs COHs + Hs$CHOHs + s CHOHs + Hs $ CH2s + H2Os
1 2 3 4 5 6 1 2 3 1 2
CO + 2s $ Cs + Os H2 + 2s $ 2Hs Cs + Hs $ CHs + s CHs + Hs $ CH2s Os + Hs $ OHs + s OHs + Hs $ H2Os + s CO + 2s $ Cs + Os H2 + s $ H2s Cs + H2s $ CH2s CO + 2s $ Cs + Os Cs + H2 $ CH2s + s
1 2 3 4
H2 + 2s $ 2Hs CO + 2Hs $ CHOHs + s CHOHs + 2Hs $ CH3OHs + 2s CH3OHs $ CH2s + H2O
1 2 3
CO + 2s $ Cs + Os H2 + 2s $ 2Hs Cs + 2Hs $ CH2s + s
FT-II
FT-III
FT-IV
FT-V
models. It can be seen that the pressure dependency of CO and H is in the range of 1/2–1, and 1/2–2, respectively. The denominator is quadratic in case of a dual-site elementary reaction, in contrast to a single-site rate-determining step. The denominator consists of the
Table 4 Reaction rate expression considered for FTS, rFT (mol g1 min1). Model
Kinetic equation
FT-I1 FT-I2
0:5 k P CO =ð1 þ aP CO þ bPH Þ
Site balance
FT-I3
0:5 0:5 =ð1 þ aPCO þ bPH Þ kPCO PH
2
s + COs + Hs s + COs + Hs
2
0:5 Þ k P H2 =ð1 þ aP CO þ bPH 2
2
s + COs + Hs
0:5 2 bPH Þ 2
s + COs + Hs
2
FT-I4 0
2
kPCO P H2 =ð1 þ aP CO þ
FT-I4
kPCO P H2 =ð1 þ
FT-I5
kPCO P H2 =ð1 þ aP CO Þ2 kPCO =ð1 þ aP CO þ bP H2 Þ kPH2 =ð1 þ aP CO þ bP H2 Þ
FT-II1 FT-II2 FT-II3
s + COs
2
kPCO P H2 =ð1 þ aP CO þ bP H2 Þ
s + COs + H2s s + COs + H2s s + COs + H2s
FT-II4
2 =ð1 þ aP CO P H2 Þ2 kPCO PH
s + CH2Os
FT-III1 FT-III2 FT-III3
kPCO =ð1 þ aP CO Þ kPCO P H2 =ð1 þ aP CO Þ
s + Cos s + Cos s + Cos
FT-III30
2 kPCO PH =ð1 2
FT-IV2
0:5 0:5 =ð1 þ aPH Þ kPCO PH
2
FT-IV3 FT-V1
2 =ð1 þ aP CO Þ kPCO PH 2
s + CH2Os
þ aP CO P H2 Þ
s + Hs
2
kPCO P H2 =ð1 þ
0:5 aPH 2
0:5 kPH2 =ð1 þ aPH Þ
þ
0:5 2 bP CO PH Þ 2
s + HOCs + Hs s + Hs
2
2
FT-V2
kPCO P H2 =ð1 þ
0:5 2 aPH Þ 2
FT-V3
2 =ð1 kPCO PH 2
0:5 aPH 2
FT-VI1 FT-VI2 FT-VII1
kPH2 =ð1 þ aP H2 Þ kPCO P H2 =ð1 þ aP H2 Þ
þ
þ
s + Hs 0:5 3 bP CO PH Þ 2
s + CH2Os + Hs s + H2s s + H2s s + Cs + Os + Hs
2
0:5 0:5 þ bPH Þ kPCO =ð1 þ aPCO 2
FT-VII2 FT-VII20
0:5 kPH2 =ð1 þ aPCO Þ
kPH2 =ð1 þ
0:5 aPCO
s + Cs + Os
2
þ
FT-IX
FT-X
individual contributions of significantly plentiful species on the catalyst surface. 3.3. Kinetic study: kinetic parameters estimation As explained in Section 3.2, twenty-two experimental tests were designed for the kinetic evaluations. The obtained results are presented in Table 1. The discrimination of the kinetic models and the estimation of the kinetic parameters were performed by fitting the experimental data of the components partial pressure to the kinetics equations (Table 4). The estimation of the values of the kinetic parameters through best-fit model (FT-IV3) was better Table 5 Parameters and mean absolute relative residuals (MARR) for the FT kinetic models.
s + HOCs
0:5 2 aP CO PH Þ 2
2
FT-VIII
s + Cs + Os + Hs
0:5 2 bPH Þ 2 0:5 2 bPH Þ 2
Model
k (x) (mol g1 min1 barx)
a (x) (barx)
b (x) (barx)
FT-I1 FT-I2 FT-I3 FT-I4 FT-I40 FT-I5 FT-II1 FT-II2 FT-II3 FT-II4 FT-III1 FT-III2 FT-III3 FT-III30 FT-IV2 FT-IV3 FT-V1 FT-V2 FT-V3 FT-VI1 FT-VI2 FT-VII1 FT-VII2 FT-VII20 FT-VII3
k1 (1) k2 (1) k3K1 (3/2) k4K1K2K3 (2) k4K1K2K3 (2) k5K1 K2 K3 K4 (2) k1 (1) k2 (1) k3 K1 K2 (2) k4 K 1 K22 (3) k1 (1) k2 (2) k3 K1 K2 (3) k3 K1 K2 (3) 1=2 k2 K1 (3/2) k3 K1 K2 (2) k1 (1) k2K1 (2) k3 K12 K 2 (3) k1 (1) k2K1 (2) k1 (1) k2 (1) k2 (1) k3(K1K2)1/2 (1)
K1 (1) K1 (1) K1 (1) K1 (1) 1=2 K 1 K2 K 3 (3/2) K1 (1) K1 (1) K1 (1) K1 (1) K1 K2 (2) K1 (1) K1 (1) K1 (1) K1 K2 (1) 1=2 K1 (1/2) 1=2 K1 (1/2) 1=2 K1 (1/2) 1=2 K1 (1/2) 1=2 K1 (1/2) K1 (1) K1 (1) 1=2 2K1 (1/2) 1=2 2K1 (1/2) 1=2 2K1 (1/2) 1=2 2K1 (1/2)
K2 1=2 K2 1=2 K2 1=2 K2
k3 K1 K 2 (3/2) k1 (1) 1=2 k2 K1 (3/2) k3 K1 K 2 (3/2)
FT-VII3
0:5 0:5 0:5 PH =ð1 þ aPCO þ kPCO
FT-VIII3
0:5 0:5 P H2 =ð1 þ aPCO þ bP H2 Þ kPCO
FT-IX1
0:5 Þ kPCO =ð1 þ aPCO
2
s + Cs + Os
FT-IX2
0:5 =ð1 kPH2 PCO
0:5 aPCO Þ
s + Cs + Os
FT-VIII3 FT-IX1 FT-IX2
FT-X3
0:5 0:5 0:5 P H2 =ð1 þ aPCO þ bPH Þ kPCO
s + Cs + Os + Hs
FT-X3
2
2
þ
3
2
s + Cs + Os + Hs s + Cs + Os + H2s
1=2
2K1 (1/2) 1=2 2K1 (1/2) 1=2 2K1 (1/2)
1=2
1=2
2K1
1=2
(1/2)
1=2
(1/2) (1/2) (1/2) (1/2)
K2 (1) K2 (1) K2 (1)
1=2
K1 K 2 (3/2)
K1K2 (3/2)
K2
1=2
(1/2)
1=2 K2 1=2 K2
(1/2) (1/2)
K2 (1)
1=2
K2
(1/2)
MARR (%) 39.47 31.10 34.48 15.83 41.11 20.70 51.75 61.01 14.95 19.09 26.27 33.29 42.74 15.10 48.29 13.03 26.18 14.67 15.49 23.56 41.90 28.40 18.92 37.69 44.49 24.68 19.22 15.93 16.16
1182
M. Mansouri et al. / Journal of Industrial and Engineering Chemistry 19 (2013) 1177–1183
determined by a multi variable non-linear regression method, using the Levenberg–Marquardt algorithm. The objective function was to minimize the sum of the square of residuals corresponding to difference between the experimental data and those calculated for the kinetic models. Arrhenius and adsorption equations were substituted in kinetics models: Eqs. (13) and (14) were substituted for the reaction rate constant (k) and the adsorption parameter group (a), respectively. E (13) k ¼ k0 exp RT
Table 6 Values of kinetic parameters of FT-IV3 model.
DH a a ¼ a0 exp RT
493.15 K. The observed rate constant for the optimal kinetic model is high in comparison to the results reported by Keyser et al. (3.84 103 mol g1 min1 bar2) [9]. It means that by decreasing the catalyst particle size, the rate constant of FTS reaction increased and as a result, FTS reaction rate increased. The calculated activation energy for the FTS reaction was found to be 80.63 kJ/mol which is close to the results reported in the literature [9–12,15,17]; the high activation energy for hydrocarbon formation suggests that the diffusion interference is not significant in the experiments [10,24]. As with intraparticle diffusion limitations, the presence of external mass-transfer limitations could be detected via measuring the apparent activation energy. An external mass-transfer control regime could lead to the apparent energy activation of just a few kJ/mol [25].
(14)
The mean absolute relative residual (MARR) between experimental and calculated consumption rate of CO is defined as: ! N exp jriexp rical j 1 X MARR ¼ (15) 100 N exp i¼1 riexp where Nexp is the number of data points included. The comparison of the calculated and experimental consumption rate of CO for the FT-IV3 model is shown in Fig. 3 and the MARR of this model was obtained 13.03%. The MARR% values of the other obtained kinetic models are presented in Table 5; as it shown the FT-IV3 model having the minimal MARR value fits the experimental data well. Model FT-IV shows that dominant mechanism on catalyst surface is based on dissociation of hydrogen with molecular CO and forming methyl monomers. Although in other studies on cobalt catalyst the dominant mechanism was carbide mechanism but it can be concluded that on bimetallic cobalt catalysts, dissociation of carbon of CO cannot be done lonely. In our previous research on titania-supported Co–Mn catalyst [19], forming of the monomer CH2 was done by reaction of adsorbed CO and hydrogen in two steps that was assumed as dominant mechanism. Also Keyser [9] observed through a study on bimetallic Co–Mn oxide catalyst that a reaction rate equation for the FT reaction based on the enolic mechanism gave results which were marginally better than results based on the carbide mechanism. The kinetic parameters and activation energy calculated for the best fitted model (FT-IV3) are presented in Table 6. The rate constant of FTS reaction over a nano-structure Co–Mn oxide catalyst was determined to be 0.65 mol g1 min1 bar2 at
Parameter k0 E a0 DHa b0 DHb
Dimension 1
Estimate 1
mol g min kJ/mol bar1/2 kJ/mol bar3/2 kJ/mol
2
bar
2.275E+08 80.63 4.438E+04 37.15 1.01E+06 62.53
4. Conclusion The nano-structured cobalt–manganese oxide catalyst was prepared by thermal decomposition of a new precursor, [Co(NH3)4CO3]MnO4. Experiments for the kinetic of Fischer– Tropsch reaction (hydrocarbon formation) were carried out over the Co–Mn oxide catalyst in a fixed bed micro-reactor over a range of operating conditions. Considering Langmuir–Hinshelwood– Hogan–Watson adsorption theory in catalytic reactions, CO consumption rate equations were defined by 10 mechanisms, consequently, 30 kinetic models were proposed. The unknown kinetic parameters were estimated from experimental data using linear regression (Levenberg–Marquardt) method. The kinetic parameters estimated for this kinetic model presented reasonable confidence intervals. The results showed that the model FT-IV3 which developed based on enolic mechanism, having the minimal MARR value fits the experimental data well. The activation energy for the best fitted model was 80.63 kJ/mol; this value of activation energy confirms that intraparticle mass transport is not significant. Kinetic results indicated that in FTS rate expression, the rate constant (k) has been increased by decreasing the catalyst particle size from conventional (prepared using co-precipitation method by Keyser et al. [9]) to nano-structured (in this study), which means that the FTS reaction rate increased by decreasing the catalyst particle size. References
Fig. 3. A comparison between the experimental data with predicted results of the FT-IV3 model equation.
[1] A.A. Adesina, Applied Catalysis A 138 (1996) 345. [2] R. Sethuraman, N.N. Bakhshi, S.P. Katikaneni, R.O. Idem, Fuel Processing Technology 73 (2001) 197. [3] A.A. Mirzaei, A. Beig babaei, M. Galavy, A. Youssefi, Fuel Processing Technology 91 (2010) 335. [4] A.A. Mirzaei, R. Habibpour, E. Kashi, Applied Catalysis A 296 (2005) 222. [5] J. Barrault, C. Forquy, V. Perrichon, Applied Catalysis A 5 (1983) 119. [6] J. Barrault, C. Forguy, J. Menezo, R. Maurel, Reaction Kinetics and Catalysis Letters 17 (1981) 373. [7] A.A. Mirzaei, M. Faizi, R. Habibpour, Applied Catalysis A 306 (2006) 98. [8] F. Morales, D. Grandjean, A. Mens, F.M.F. de Groot, B.M. Weckhuysen, Journal of Physical Chemistry B 110 (2006) 8626. [9] M.J. Keyser, R.C. Everson, R.L. Espinoza, Industrial and Engineering Chemistry Research 39 (2000) 48. [10] A. Sari, Y. Zamani, S.A. Taheri, Fuel Processing Technology 90 (2009) 1305. [11] I.C. Yates, C.N. Satterfield, Energy and Fuels 5 (1991) 168.
M. Mansouri et al. / Journal of Industrial and Engineering Chemistry 19 (2013) 1177–1183 [12] B. Sarup, B.W. Wojciechowski, Canadian Journal of Chemical Engineering 67 (1989) 62. [13] M. van der Reit, Carbon Monoxide Hydrogenation with Transition Metal Oxide Supported Cobalt and Iron Catalysts, Ph.D. Thesis, University of the Witwatersrand, Johanssesburg, South Africa, 1998. [14] J. Yang, Y. Liu, J. Chang, Y.N. Wang, L. Bai, Y.Y. Xu, H.W. Xiang, Y.W. Li, B. Zhang, Industrial and Engineering Chemistry Research 42 (2003) 5066. [15] C.G. Visconti, E. Tronconi, L. Lietti, R. Zennaro, P. Forzatti, Chemical Engineering Science 62 (2007) 5338. [16] B.T. Teng, J. Chang, C.H. Zhang, D.B. Cao, J. Yang, Y. Liu, X.H. Guo, H.W. Xiang, Y.W. Li, Applied Catalysis A 301 (2006) 39. ¨ nsan, Journal of Chemical Technology and Biotechnology 70 [17] A.N. Akın, Z.I˙. O (1997) 304.
1183
[18] H.H. Storch, N. Golumbic, R.B. Aderson, The Fischer–Tropsch and Related Synthesis, Wiley, New York, 1951. [19] H. Atashi, F. Siami, A.A. Mirzaei, M. Sarkari, Journal of Industrial and Engineering Chemistry 16 (2010) 952. [20] H. Atashi, M. Mansouri, S.H. Hosseini, M. Khorram, A.A. Mirzaei, M. Karimi, G. Mansouri, Korean Journal of Chemical Engineering 29 (2012) 304. [21] A.A. Mirzaei, B. Shirzadi, H. Atashi, M. Mansouri, Journal of Industrial and Engineering Chemistry 18 (2012) 1515. [22] B.H. Davis, Fuel Processing Technology 71 (2001) 157. [23] H. Chen, A.A. Adesina, Applied Catalysis A 112 (1994) 87. [24] J. Chang, L. Bai, B.T. Teng, R. Zhang, J. Yang, Y.Y. Xu, H.W. X iang, Y.W. Li, Chemical Engineering Science 62 (2007) 4983. [25] A.Y. Khodakov, W. Chu, P. Fongarland, Chemical Reviews 107 (2007) 1692.
Contents lists available at SciVerse ScienceDirect
Journal of Industrial and Engineering Chemistry journal homepage: www.elsevier.com/locate/jiec
Kinetics studies of nano-structured cobalt–manganese oxide catalysts in Fischer–Tropsch synthesis Mohsen Mansouri a, Hossein Atashi a,*, Farshad Farshchi Tabrizi a, Ali Akbar Mirzaei b, Ghobad mansouri c a
Department of Chemical Engineering, University of Sistan & Baluchestan, Zahedan 98164-161, Iran Department of Chemistry, University of Sistan & Baluchestan, Zahedan 98135-674, Iran c Department of Chemistry, University of Payame Noor, Ilam 69319-36173, Iran b
A R T I C L E I N F O
Article history: Received 24 September 2012 Accepted 11 December 2012 Available online 20 December 2012 Keywords: [Co(NH3)4CO3]MnO4 precursor Cobalt–manganese catalyst Fischer–Tropsch synthesis Kinetic model
A B S T R A C T
The nano-structured cobalt/manganese oxide catalyst was prepared by thermal decomposition of [Co(NH3)4CO3]MnO4 precursor, and was tested for the Fischer–Tropsch reaction (hydrocarbon forming) in a fixed-bed micro-reactor. Experimental conditions were varied as follow: reaction pressure 1–10 bar, H2/CO feed ratio of 1–2 and space velocity of 3600 h1 at the temperature range of 463.15–523.15 K. On the basis of carbide and/or enolic mechanisms and Langmuir–Hinshelwood–Hougen–Watson (LHHW) type rate equations, 30 kinetic expressions for CO consumption were tested and interaction between adsorption HCO and dissociated adsorption hydrogen as the controlling step gave the most plausible kinetic model. The kinetic parameters were estimated with non-linear regression method and the activation energy was 80.63 kJ/mol for optimal kinetic model. Kinetic results indicated that in Fischer– Tropsch synthesis (FTS) rate expression, the rate constant (k) has been increased by decreasing the catalyst particle size. The catalyst characterization was carried out using different methods including powder X-ray diffraction (XRD), scanning electron microscopy (SEM) and Brunauer–Emmett–Teller (BET) surface area measurements. ß 2012 The Korean Society of Industrial and Engineering Chemistry. Published by Elsevier B.V. All rights reserved.
1. Introduction Nowadays, there is a renewed interest in Fischer–Tropsch synthesis (FTS), especially for the selective production of petrochemical feedstock such as ethylene, propylene and butylene (C2–C4 olefins) directly from synthesis gas [1,2]. Due to the thermodynamic and kinetic limitations of the reaction, it is believed that bimetallic catalysts system containing one or two metal active agents increase the value of the light olefins whereas catalyst composition exerts the greatest influence on the molecular weight distribution for the FTS products [3,4]. It is shown that the addition of Mn to Fe or Co catalysts leads to a significant increase in high olefin formation and a decrease in methane activity [5,6]. Co–Mn catalysts have been investigated intensively for its higher selectivity to lower molecular weight olefins [7,8]. The kinetic description of the FT reaction is extremely important for the industrial practice, being a prerequisite for the industrial FT process design, scale-up, optimization, and simulation. Intrinsic reaction rate equations for the FT equation
* Corresponding author. E-mail address: [email protected] (H. Atashi).
(hydrocarbon forming only) based on Langmuir–Hinshelwood– Hougen–Watson (LHHW) adsorption theory have been developed and used successfully for iron- and cobalt-based catalysts [9–21]. The mechanisms proposed include the carbide, enolic, and direct insertion theories [22]. An investigation was undertaken to elucidate the reaction kinetics of the Co–Mn oxide catalysts used by Keyser et al. [9] and to compare its performance with other cobalt- and iron-based catalysts. It was found that a reaction rate equation for FT reaction based on the enolic theory performed slightly better than a reaction rate equation based on the carbide theory. Chen and Adesina [23] have studied improved alkene selectivity over a silica supported cobalt–molybdenum catalyst, and their data also suggest the carbide mechanism where the rate determining step is the dissociative adsorption of hydrogen. Sari et al. [10] studied the rate of FTS over alumina-supported Co–Ru in a slurry phase reactor. They considered five parametric kinetic expressions based on their previous studies: one empirical power law model and four variations of the LHHW approach regarding strong inhibition due to CO adsorption and the kinetic data, fitted by a simple LHHW. Although numerous kinetic expressions were developed for Co-based catalyst, one should consider the kinetics of bimetallic
1226-086X/$ – see front matter ß 2012 The Korean Society of Industrial and Engineering Chemistry. Published by Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.jiec.2012.12.015
M. Mansouri et al. / Journal of Industrial and Engineering Chemistry 19 (2013) 1177–1183
1178
2.2. Catalyst characterization Nomenclature r FT a b k K PCO P H2 0 FCO R T W X r CO s
uS uHOC uH E
DH N
(mol g1 cat
1
rate of reaction min ) adsorption parameter (see Table 4) adsorption parameter (see Table 4) reaction rate constant (see Table 4) adsorption constant (see Table 4) CO pressure (bar) hydrogen pressure (bar) molar flow rate of CO at the inlet (mol min1) gas constant (8.314 J mol1 K1) temperature (K) catalyst mass (g) conversion the consumption rate of CO unoccupied active sites the concentration of free sites surface occupied with HOC surface occupied with H activation energy (kJ/mol) heat of adsorption (kJ/mol) number experimental data points
Superscripts experimental value exp conditions in feed in predicted value cal
Co-based FT catalyst. Thus, the main objective of the present work is to investigate the kinetic and mechanism of the CO hydrogenation on nano-structured Co–Mn catalyst, which was prepared by thermal decomposition of [Co(NH3)4CO3]MnO4 precursor. We establish kinetic models for hydrocarbon formation on the basis of the correlation between experimental data and supposed reaction mechanism sets. The suitable model was obtained and the kinetic parameters were appointed. The catalyst characterization was carried out using different methods including XRD, SEM and BET. 2. Experimental
The catalyst was analyzed by X-ray diffraction and all peaks were consistent with the peaks standard Co–Mn oxide. No peaks from any other phases of Co–Mn oxide were observed. The XRD pattern when compared with JCPDS reveals the structure as spinel oxide and a mixed oxide spinel structure (CoMn) (CoMn)2O4 was identified which is similar to that reported by van der Reit [13]. The morphology of catalysts and their precursors was observed by means of an S-360 scanning electron microscopy (USA). BET surface areas, pore volumes and average pore sizes of the catalyst precursor and calcined samples were measured by N2 physisorption using a Quantachrome Nova 2000 automated system (USA). Each catalyst sample was degassed under nitrogen atmosphere at 300 8C for 3 h. In order to obtain the BET surface areas, pore volumes and average pore sizes, different samples were evacuated at 196 8C for 66 min. 2.3. Catalyst testing The kinetic experiments were carried out in a tubular fixed bed micro-reactor. A schematic representation of the reactor is shown in Fig. 1. The reactor consisted of a single stainless steel tube with an inner diameter of 20 mm surrounded by an alumina jacket to achieve a uniform wall temperature along the length of the reactor. A preheating zone a head of catalyst packing was filled with inert quartz glass beads. External heating was provided by an electrical element wrapped around the alumina jacket and placed in through the fire brick part. The required amount of the catalyst is diluted by asbestos and placed among the inert quartz glass beads. The temperatures of all of different zones include: preheating zone, catalyst bed and underneath zone of the reactor is checked by three separate thermocouples placed in different parts of the reactor. The thermocouple which controls bed temperature is placed exactly below the catalyst bed. The inlet feed gas is arrived from top of the reactor and the outlet products are avoided from underside of the reactor. All gas lines to the reactor bed were made from 1/400 stainless steel tubing. Three mass flow controllers (Brooks, Model, 5850E) were used to adjust automatically flow rate of the inlet gases comprising CO, H2 and N2 (purity of 99.99%). Mixture of CO, H2 and N2 were subsequently introduced into the reactor, which was placed inside a tubular furnace (Atbin, Model ATU 150-15). Prior to the reaction the catalysts were reduced in situ using H2 (30 ml/min) and N2 (30 ml/min) mixture gas at 350 8C for 16 h. In
2.1. Catalyst preparation and characterization The catalyst used in the present study was prepared using a thermal decomposition precursor as follows. An aqueous solution (60 ml) of ammonium carbonate (20 g) was added to a concentrated aqueous NH3 (60 ml), and the mixture was stirred at room temperature for 30 min. To this solution was then added an aqueous solution (30 ml) of Co(NH3)26H2O (15 g) and the resulting solution was stirred at room temperature for 1 h. Then, 8 ml of 30% hydrogen peroxide was dropwise added into the solution, stirring continuously. The final solution was filtered and left for slow evaporation in air until water-soluble [Co(NH3)4CO3]NO3 crystals were obtained. To a solution containing [Co(NH3)4CO3]NO3 (1 mmol) in 100 ml water was added KMnO4 (1 mmol), and the mixture was stirred for several minutes. The precipitates ([Co(NH3)4CO3]MnO4 precursor) were recovered by centrifugation and washed by distilled water and then dried in ambient air at 30 8C overnight. In order to obtain the final nanocatalyst, the precursor was calcined at 400 8C in static air in the electric furnace for 4 h. The gray powder, nano-structured Co–Mn oxide catalyst was formed and kept in desiccator.
Fig. 1. Different parts of micro-fixed-bed reactor.
M. Mansouri et al. / Journal of Industrial and Engineering Chemistry 19 (2013) 1177–1183
each test, 1.0 g catalyst was loaded and the reactor operated about 12 h to ensure attaining the steady state operating conditions. The catalysts were extremely fine particles so intraparticle diffusion could be neglected. The gas hourly space velocity (GHSV) increased to the value in which the CO conversion was almost the same for a variety of catalyst weight which indicates that external diffusion can be neglected above this GHSV. Hence, the kinetic experiments were conducted free from internal and external mass transfer limitations. Experiments were carried out with mixtures of H2, CO and N2 in a temperature range of 463.15–523.15 K, H2/CO feed ratio of 1/1–2/1, pressure range of 1–10 bar and GHSV = 3600 h1. The catalytic test data are represented in Table 1. The stability of the catalyst was investigated by repetition of central points of the designs in the middle and the end of the experiments. We have a differential flow reactor when we choose to consider the rate to be constant at all points within the reactor. Since rates are concentration-dependent this assumption is usually reasonable only for small conversions or for shallow small reactors. For conversions up to 15%, the curve conversion versus inverse rate is nearly constant; hence the assumption of differential reactor can be used. For each run in a differential reactor, the plug flow performance equation becomes as follows: W ¼ 0 FCO
Z
xout xin
dx xout xin xout 0 ¼ ¼ r CO r CO r CO
(1)
According to the above equation, the average rate for each run is derived as follow: r CO ¼
1179
several agglomerates of irregularly spherical shaped grains and shows that this material has a less dense and homogeneous morphology. After the calcination at 400 8C for 4 h, the morphological features are different with the precursor sample and show that the agglomerate size is greatly reduced compared to the precursor (Fig. 2b). It may be due to that the calcined catalyst surface is covered with small crystallite of cobalt and manganese oxide. After FTS chemical reaction the catalyst texture and its morphology changed (Fig. 2c). However, the size of these grains grew larger by agglomeration in the tested catalyst, which may be due to the sintering after reactions. 3.1.2. BET measurements BET surface area, pore volume and pore diameter of the precursor, nano and conventional (prepared using co-precipitation method by Keyser et al. [9]) catalysts are summarized in Table 2. As shown from this Table, the BET surface area and pore volume of precursor are low, about 11.32 m2/g and 0.12 cm3/g, respectively. It can be seen that after calcinations, both the BET surface area and pore volume increased to 74.28 m2/g and 0.21 cm3/g, respectively; this is in agreement with the SEM results which showed that the agglomerate size of calcined catalyst was less compared with its precursor and therefore leaded to an increase in the BET specific surface area of the calcined sample. Also, it can be found that the BET surface area of the nano catalyst is much higher than the conventional catalyst used by Keyser et al. [9]. 3.2. Kinetic modeling and reaction rate equations
0 FCO xout W
(2)
3. Result and discussion 3.1. Catalyst characterization 3.1.1. Morphological properties The scanning electron microscopy (SEM) observations have shown differences in morphology of precursor and calcined catalysts (before and after the reaction) and indicate that these materials are made of nano-metric particles. The electron micrograph obtained from catalyst precursor (Fig. 2a) depicts
Although great efforts have been devoted to elucidate the FTS reaction mechanism, there are still many controversies on this point, and many theories have been proposed in the literature [14– 16]. On the basis of the nature of CO adsorption and of the nature of chain initiator intermediates, popular mechanistic proposals include (1) the carbide mechanism [17,22], wherein CO adsorbs dissociatively and the carbide (Cs) is the chain initiator intermediate. In this mechanism simultaneous dissociative adsorption of CO and H2 is followed by the hydrogenation of adsorbed carbon by adsorbed hydrogen in a stepwise manner to give methane and higher hydrocarbons. (2) The enolic mechanism [18,22], involving molecular adsorption of CO which reacts with adsorbed hydrogen
Table 1 Summary of experimental conditions and results for kinetic tests at PTot. = 1–10 bar, T = 463.15–523.15 K and GHSV = 3600 h1 in a fixed bed reactor (FBR). No.
T (K)
H2/CO
P CO (bar)
P H2 (bar)
0 FCO 103 (mol min1)
XCO (%)
1 r CO 102 (mol g1 ) cat min
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22
463.15 463.15 473.15 488.15 498.15 498.15 498.15 463.15 463.15 488.15 488.15 493.15 493.15 498.15 498.15 463.15 463.15 468.15 488.15 498.15 503.15 523.15
1 1 1 1 1 1 1 1.5 1.5 1.5 1.5 1.5 1.5 1.5 1.5 2 2 2 2 2 2 2
0.9 2.1 1.5 0.9 0.9 1.5 2.1 0.72 1.2 1.2 1.68 0.72 1.2 1.68 0.72 1.4 2 0.6 2 0.6 1.4 2
0.9 2.1 1.5 0.9 0.9 1.5 2.1 1.08 1.8 1.8 2.52 1.08 1.8 2.52 1.08 2.8 4 1.2 4 1.2 2.8 4
1.402 3.258 2.288 1.330 1.304 2.173 3.042 1.119 1.858 1.774 2.484 1.058 1.745 2.434 1.043 2.181 3.116 0.925 2.957 0.867 2.008 2.759
8.448 9.28 7.7 8.58 9.12 9.84 11.82 9.47 9.69 12.48 11.79 12.75 12.36 11.91 13.28 9.615 9.762 8.93 11.39 12.5 12.17 14.21
1.185 3.023 1.762 1.142 1.188 2.138 3.595 1.060 1.800 2.214 2.928 1.349 2.157 2.899 1.385 2.097 3.042 0.871 3.369 1.143 2.444 3.920
M. Mansouri et al. / Journal of Industrial and Engineering Chemistry 19 (2013) 1177–1183
1180
Fig. 2. SEM micrographs of: (a) precursor, (b) calcined catalyst before the test, and (c) calcined catalyst after the test.
and the formation of an oxygenated intermediate, the enol (HOCs). In this mechanism, dissociative adsorption of H2 and molecular adsorption of CO followed by the hydrogenation of adsorbed carbon monoxide by adsorbed hydrogen to form an oxygenated intermediate which reacts with another adsorbed hydrogen to form water and adsorbed carbon, and the reaction of the resulting carbon with adsorbed hydrogen as in the carbide mechanism. To obtain the rate equation, a reaction mechanism should be adopted. Based on the elementary reactions explained in the above, we defined ten different possible mechanisms. The elementary reactions, which have been settled on sites for different models are listed in Table 3. These mechanisms were offered on the basis of various monomer and carbon chain distribution pathway. These mechanisms were obtained by considering different possibility of adsorption of CO and H2 molecules (associative or dissociative mode) on the catalyst surface. Each of these adsorption modes leads to different mechanisms and hence different kinetic equations were obtained. In order to derive rate equations, the LHHW approach was utilized and for each model, the possible ratedetermining steps were identified, while all other steps were assumed to be at quasi-equilibrium. Then, all of the obtained models were fitted separately, against experimental data. Table 4 displays the final form of different rate expressions for the 10 kinetic models. The development of the kinetic equations will be illustrated for model FT-IV3 (the third step elementary reaction from the FT-IV model). The model codes refer to the set of elementary reactions and the elementary reaction is not at equilibrium (that is the rate-determining step, so in this case reaction (3)). The set of elementary reactions for model FT-IV3 is shown in Table 3. Firstly, dissociated hydrogen reacted with molecular CO yielding formyl (HOC) intermediates (step 2) and these species were hydrogenated in steps 3 and 4, forming methyl monomers. The reaction rate of the rate-determining step is: r FT-IV3 ¼ k3 uHOC uH
(3)
where uHOC is the surface fraction occupied with the formyl intermediate and uH is the surface fraction occupied with the dissociative adsorbed hydrogen. The fraction of vacant sites, uS, can be calculated from the following balance equation:
quasi-equilibrium: k1
H2 þ 2s !2Hs 2
(5)
2
k1 P H2 uS k1 uH ¼ 0
uH ¼ K10:5 PH0:52 uS K1 ¼
(6)
k1 k1
where K1 is the equilibrium constant of dissociated hydrogen adsorption step. k2
CO þ Hs !HOCs
(7)
k2 P CO uH k2 uHOC ¼ 0
uHOC ¼ K 2 PCO uH ¼ K10:5 K 2 PH0:52 PCO uS K2 ¼
(8)
k2 k2
Substituting Eqs. (6) and (8) into Eq. (4), the concentration of free active site can be expressed as:
uS ¼
1 0:5 þ K P 0:5 1 þ K10:5 PH 2 CO PH2 2
(9)
with substituting of Eq. (9) into Eqs. (6) and (8), the expressions of [uH] and [uHCO] becomes:
uH ¼
1þ
uHOC ¼
0:5 K10:5 PH 2 0:5 0:5 0:5 K1 PH2 þ K 2 PCO PH 2
(10)
0:5 K10:5 K 2 PH P CO 2
(11)
0:5 þ K P 0:5 1 þ K10:5 PH 2 CO PH2 2
By substituting of the surface fraction of HOC and H in Eq. (3), the final rate expression is obtained as follows:
r FT-IV3 ¼
k3 K 1 K 2 PCO P H2 0:5 þ K P 0:5 ð1 þ K10:5 PH 2 CO PH2 Þ 2
2
¼
kPCO P H2 2
0:5 þ bP 0:5 ð1 þ aPH CO PH2 Þ 2
(12)
uS þ uH þ uHOC ¼ 1
(4)
The surface fractions of HOC and H can be calculated from the site balance, the preceding reaction steps which are at
Table 4 summarizes the final form of the different rate expressions for the 30 possible kinetic models, whereas Table 5 shows the kinetic and adsorption parameters for the several kinetic
Table 2 Textural of the precursor, nano and conventional catalysts. Sample
BET surface area (m2/g)
Pore volume (cm3/g)
Average pore diameter (nm)
[Co(NH3)4CO3]MnO4 Nano catalyst Conventional catalyst [9]
11.32 74.28 4.3
0.12 0.21 NA
16.89 13.32 NA
M. Mansouri et al. / Journal of Industrial and Engineering Chemistry 19 (2013) 1177–1183
1181
Table 3 Elementary reactions sets for FT synthesis. Model
Number
Elementary reaction
Model
Number
Elementary reaction
FT-I
1 2 3 4 5
CO + s $ COs H2 + 2s $ 2Hs COs + Hs $ CHOs + s CHOs + Hs $ CHOHs + s CHOHs + s $ CHs + OHs
FT-VI
1 2 3
H2 + S $ H2s CO + H2s $ CHOHs CHOHs + H2s $ CH2s + H2Os
FT-VII
1 2 3 4 1 2 3 1 2 3 4
CO + s $ COs H2 + s $ H2s COs + H2s+ $ CH2Os + s CH2Os+H2s $ CH2s+H2O CO + s $ COs COs + H2 $ CHOHs CHOHs + H2 $ CH2s + H2O H2 + 2s $ 2Hs CO + Hs $ COHs COHs + Hs$CHOHs + s CHOHs + Hs $ CH2s + H2Os
1 2 3 4 5 6 1 2 3 1 2
CO + 2s $ Cs + Os H2 + 2s $ 2Hs Cs + Hs $ CHs + s CHs + Hs $ CH2s Os + Hs $ OHs + s OHs + Hs $ H2Os + s CO + 2s $ Cs + Os H2 + s $ H2s Cs + H2s $ CH2s CO + 2s $ Cs + Os Cs + H2 $ CH2s + s
1 2 3 4
H2 + 2s $ 2Hs CO + 2Hs $ CHOHs + s CHOHs + 2Hs $ CH3OHs + 2s CH3OHs $ CH2s + H2O
1 2 3
CO + 2s $ Cs + Os H2 + 2s $ 2Hs Cs + 2Hs $ CH2s + s
FT-II
FT-III
FT-IV
FT-V
models. It can be seen that the pressure dependency of CO and H is in the range of 1/2–1, and 1/2–2, respectively. The denominator is quadratic in case of a dual-site elementary reaction, in contrast to a single-site rate-determining step. The denominator consists of the
Table 4 Reaction rate expression considered for FTS, rFT (mol g1 min1). Model
Kinetic equation
FT-I1 FT-I2
0:5 k P CO =ð1 þ aP CO þ bPH Þ
Site balance
FT-I3
0:5 0:5 =ð1 þ aPCO þ bPH Þ kPCO PH
2
s + COs + Hs s + COs + Hs
2
0:5 Þ k P H2 =ð1 þ aP CO þ bPH 2
2
s + COs + Hs
0:5 2 bPH Þ 2
s + COs + Hs
2
FT-I4 0
2
kPCO P H2 =ð1 þ aP CO þ
FT-I4
kPCO P H2 =ð1 þ
FT-I5
kPCO P H2 =ð1 þ aP CO Þ2 kPCO =ð1 þ aP CO þ bP H2 Þ kPH2 =ð1 þ aP CO þ bP H2 Þ
FT-II1 FT-II2 FT-II3
s + COs
2
kPCO P H2 =ð1 þ aP CO þ bP H2 Þ
s + COs + H2s s + COs + H2s s + COs + H2s
FT-II4
2 =ð1 þ aP CO P H2 Þ2 kPCO PH
s + CH2Os
FT-III1 FT-III2 FT-III3
kPCO =ð1 þ aP CO Þ kPCO P H2 =ð1 þ aP CO Þ
s + Cos s + Cos s + Cos
FT-III30
2 kPCO PH =ð1 2
FT-IV2
0:5 0:5 =ð1 þ aPH Þ kPCO PH
2
FT-IV3 FT-V1
2 =ð1 þ aP CO Þ kPCO PH 2
s + CH2Os
þ aP CO P H2 Þ
s + Hs
2
kPCO P H2 =ð1 þ
0:5 aPH 2
0:5 kPH2 =ð1 þ aPH Þ
þ
0:5 2 bP CO PH Þ 2
s + HOCs + Hs s + Hs
2
2
FT-V2
kPCO P H2 =ð1 þ
0:5 2 aPH Þ 2
FT-V3
2 =ð1 kPCO PH 2
0:5 aPH 2
FT-VI1 FT-VI2 FT-VII1
kPH2 =ð1 þ aP H2 Þ kPCO P H2 =ð1 þ aP H2 Þ
þ
þ
s + Hs 0:5 3 bP CO PH Þ 2
s + CH2Os + Hs s + H2s s + H2s s + Cs + Os + Hs
2
0:5 0:5 þ bPH Þ kPCO =ð1 þ aPCO 2
FT-VII2 FT-VII20
0:5 kPH2 =ð1 þ aPCO Þ
kPH2 =ð1 þ
0:5 aPCO
s + Cs + Os
2
þ
FT-IX
FT-X
individual contributions of significantly plentiful species on the catalyst surface. 3.3. Kinetic study: kinetic parameters estimation As explained in Section 3.2, twenty-two experimental tests were designed for the kinetic evaluations. The obtained results are presented in Table 1. The discrimination of the kinetic models and the estimation of the kinetic parameters were performed by fitting the experimental data of the components partial pressure to the kinetics equations (Table 4). The estimation of the values of the kinetic parameters through best-fit model (FT-IV3) was better Table 5 Parameters and mean absolute relative residuals (MARR) for the FT kinetic models.
s + HOCs
0:5 2 aP CO PH Þ 2
2
FT-VIII
s + Cs + Os + Hs
0:5 2 bPH Þ 2 0:5 2 bPH Þ 2
Model
k (x) (mol g1 min1 barx)
a (x) (barx)
b (x) (barx)
FT-I1 FT-I2 FT-I3 FT-I4 FT-I40 FT-I5 FT-II1 FT-II2 FT-II3 FT-II4 FT-III1 FT-III2 FT-III3 FT-III30 FT-IV2 FT-IV3 FT-V1 FT-V2 FT-V3 FT-VI1 FT-VI2 FT-VII1 FT-VII2 FT-VII20 FT-VII3
k1 (1) k2 (1) k3K1 (3/2) k4K1K2K3 (2) k4K1K2K3 (2) k5K1 K2 K3 K4 (2) k1 (1) k2 (1) k3 K1 K2 (2) k4 K 1 K22 (3) k1 (1) k2 (2) k3 K1 K2 (3) k3 K1 K2 (3) 1=2 k2 K1 (3/2) k3 K1 K2 (2) k1 (1) k2K1 (2) k3 K12 K 2 (3) k1 (1) k2K1 (2) k1 (1) k2 (1) k2 (1) k3(K1K2)1/2 (1)
K1 (1) K1 (1) K1 (1) K1 (1) 1=2 K 1 K2 K 3 (3/2) K1 (1) K1 (1) K1 (1) K1 (1) K1 K2 (2) K1 (1) K1 (1) K1 (1) K1 K2 (1) 1=2 K1 (1/2) 1=2 K1 (1/2) 1=2 K1 (1/2) 1=2 K1 (1/2) 1=2 K1 (1/2) K1 (1) K1 (1) 1=2 2K1 (1/2) 1=2 2K1 (1/2) 1=2 2K1 (1/2) 1=2 2K1 (1/2)
K2 1=2 K2 1=2 K2 1=2 K2
k3 K1 K 2 (3/2) k1 (1) 1=2 k2 K1 (3/2) k3 K1 K 2 (3/2)
FT-VII3
0:5 0:5 0:5 PH =ð1 þ aPCO þ kPCO
FT-VIII3
0:5 0:5 P H2 =ð1 þ aPCO þ bP H2 Þ kPCO
FT-IX1
0:5 Þ kPCO =ð1 þ aPCO
2
s + Cs + Os
FT-IX2
0:5 =ð1 kPH2 PCO
0:5 aPCO Þ
s + Cs + Os
FT-VIII3 FT-IX1 FT-IX2
FT-X3
0:5 0:5 0:5 P H2 =ð1 þ aPCO þ bPH Þ kPCO
s + Cs + Os + Hs
FT-X3
2
2
þ
3
2
s + Cs + Os + Hs s + Cs + Os + H2s
1=2
2K1 (1/2) 1=2 2K1 (1/2) 1=2 2K1 (1/2)
1=2
1=2
2K1
1=2
(1/2)
1=2
(1/2) (1/2) (1/2) (1/2)
K2 (1) K2 (1) K2 (1)
1=2
K1 K 2 (3/2)
K1K2 (3/2)
K2
1=2
(1/2)
1=2 K2 1=2 K2
(1/2) (1/2)
K2 (1)
1=2
K2
(1/2)
MARR (%) 39.47 31.10 34.48 15.83 41.11 20.70 51.75 61.01 14.95 19.09 26.27 33.29 42.74 15.10 48.29 13.03 26.18 14.67 15.49 23.56 41.90 28.40 18.92 37.69 44.49 24.68 19.22 15.93 16.16
1182
M. Mansouri et al. / Journal of Industrial and Engineering Chemistry 19 (2013) 1177–1183
determined by a multi variable non-linear regression method, using the Levenberg–Marquardt algorithm. The objective function was to minimize the sum of the square of residuals corresponding to difference between the experimental data and those calculated for the kinetic models. Arrhenius and adsorption equations were substituted in kinetics models: Eqs. (13) and (14) were substituted for the reaction rate constant (k) and the adsorption parameter group (a), respectively. E (13) k ¼ k0 exp RT
Table 6 Values of kinetic parameters of FT-IV3 model.
DH a a ¼ a0 exp RT
493.15 K. The observed rate constant for the optimal kinetic model is high in comparison to the results reported by Keyser et al. (3.84 103 mol g1 min1 bar2) [9]. It means that by decreasing the catalyst particle size, the rate constant of FTS reaction increased and as a result, FTS reaction rate increased. The calculated activation energy for the FTS reaction was found to be 80.63 kJ/mol which is close to the results reported in the literature [9–12,15,17]; the high activation energy for hydrocarbon formation suggests that the diffusion interference is not significant in the experiments [10,24]. As with intraparticle diffusion limitations, the presence of external mass-transfer limitations could be detected via measuring the apparent activation energy. An external mass-transfer control regime could lead to the apparent energy activation of just a few kJ/mol [25].
(14)
The mean absolute relative residual (MARR) between experimental and calculated consumption rate of CO is defined as: ! N exp jriexp rical j 1 X MARR ¼ (15) 100 N exp i¼1 riexp where Nexp is the number of data points included. The comparison of the calculated and experimental consumption rate of CO for the FT-IV3 model is shown in Fig. 3 and the MARR of this model was obtained 13.03%. The MARR% values of the other obtained kinetic models are presented in Table 5; as it shown the FT-IV3 model having the minimal MARR value fits the experimental data well. Model FT-IV shows that dominant mechanism on catalyst surface is based on dissociation of hydrogen with molecular CO and forming methyl monomers. Although in other studies on cobalt catalyst the dominant mechanism was carbide mechanism but it can be concluded that on bimetallic cobalt catalysts, dissociation of carbon of CO cannot be done lonely. In our previous research on titania-supported Co–Mn catalyst [19], forming of the monomer CH2 was done by reaction of adsorbed CO and hydrogen in two steps that was assumed as dominant mechanism. Also Keyser [9] observed through a study on bimetallic Co–Mn oxide catalyst that a reaction rate equation for the FT reaction based on the enolic mechanism gave results which were marginally better than results based on the carbide mechanism. The kinetic parameters and activation energy calculated for the best fitted model (FT-IV3) are presented in Table 6. The rate constant of FTS reaction over a nano-structure Co–Mn oxide catalyst was determined to be 0.65 mol g1 min1 bar2 at
Parameter k0 E a0 DHa b0 DHb
Dimension 1
Estimate 1
mol g min kJ/mol bar1/2 kJ/mol bar3/2 kJ/mol
2
bar
2.275E+08 80.63 4.438E+04 37.15 1.01E+06 62.53
4. Conclusion The nano-structured cobalt–manganese oxide catalyst was prepared by thermal decomposition of a new precursor, [Co(NH3)4CO3]MnO4. Experiments for the kinetic of Fischer– Tropsch reaction (hydrocarbon formation) were carried out over the Co–Mn oxide catalyst in a fixed bed micro-reactor over a range of operating conditions. Considering Langmuir–Hinshelwood– Hogan–Watson adsorption theory in catalytic reactions, CO consumption rate equations were defined by 10 mechanisms, consequently, 30 kinetic models were proposed. The unknown kinetic parameters were estimated from experimental data using linear regression (Levenberg–Marquardt) method. The kinetic parameters estimated for this kinetic model presented reasonable confidence intervals. The results showed that the model FT-IV3 which developed based on enolic mechanism, having the minimal MARR value fits the experimental data well. The activation energy for the best fitted model was 80.63 kJ/mol; this value of activation energy confirms that intraparticle mass transport is not significant. Kinetic results indicated that in FTS rate expression, the rate constant (k) has been increased by decreasing the catalyst particle size from conventional (prepared using co-precipitation method by Keyser et al. [9]) to nano-structured (in this study), which means that the FTS reaction rate increased by decreasing the catalyst particle size. References
Fig. 3. A comparison between the experimental data with predicted results of the FT-IV3 model equation.
[1] A.A. Adesina, Applied Catalysis A 138 (1996) 345. [2] R. Sethuraman, N.N. Bakhshi, S.P. Katikaneni, R.O. Idem, Fuel Processing Technology 73 (2001) 197. [3] A.A. Mirzaei, A. Beig babaei, M. Galavy, A. Youssefi, Fuel Processing Technology 91 (2010) 335. [4] A.A. Mirzaei, R. Habibpour, E. Kashi, Applied Catalysis A 296 (2005) 222. [5] J. Barrault, C. Forquy, V. Perrichon, Applied Catalysis A 5 (1983) 119. [6] J. Barrault, C. Forguy, J. Menezo, R. Maurel, Reaction Kinetics and Catalysis Letters 17 (1981) 373. [7] A.A. Mirzaei, M. Faizi, R. Habibpour, Applied Catalysis A 306 (2006) 98. [8] F. Morales, D. Grandjean, A. Mens, F.M.F. de Groot, B.M. Weckhuysen, Journal of Physical Chemistry B 110 (2006) 8626. [9] M.J. Keyser, R.C. Everson, R.L. Espinoza, Industrial and Engineering Chemistry Research 39 (2000) 48. [10] A. Sari, Y. Zamani, S.A. Taheri, Fuel Processing Technology 90 (2009) 1305. [11] I.C. Yates, C.N. Satterfield, Energy and Fuels 5 (1991) 168.
M. Mansouri et al. / Journal of Industrial and Engineering Chemistry 19 (2013) 1177–1183 [12] B. Sarup, B.W. Wojciechowski, Canadian Journal of Chemical Engineering 67 (1989) 62. [13] M. van der Reit, Carbon Monoxide Hydrogenation with Transition Metal Oxide Supported Cobalt and Iron Catalysts, Ph.D. Thesis, University of the Witwatersrand, Johanssesburg, South Africa, 1998. [14] J. Yang, Y. Liu, J. Chang, Y.N. Wang, L. Bai, Y.Y. Xu, H.W. Xiang, Y.W. Li, B. Zhang, Industrial and Engineering Chemistry Research 42 (2003) 5066. [15] C.G. Visconti, E. Tronconi, L. Lietti, R. Zennaro, P. Forzatti, Chemical Engineering Science 62 (2007) 5338. [16] B.T. Teng, J. Chang, C.H. Zhang, D.B. Cao, J. Yang, Y. Liu, X.H. Guo, H.W. Xiang, Y.W. Li, Applied Catalysis A 301 (2006) 39. ¨ nsan, Journal of Chemical Technology and Biotechnology 70 [17] A.N. Akın, Z.I˙. O (1997) 304.
1183
[18] H.H. Storch, N. Golumbic, R.B. Aderson, The Fischer–Tropsch and Related Synthesis, Wiley, New York, 1951. [19] H. Atashi, F. Siami, A.A. Mirzaei, M. Sarkari, Journal of Industrial and Engineering Chemistry 16 (2010) 952. [20] H. Atashi, M. Mansouri, S.H. Hosseini, M. Khorram, A.A. Mirzaei, M. Karimi, G. Mansouri, Korean Journal of Chemical Engineering 29 (2012) 304. [21] A.A. Mirzaei, B. Shirzadi, H. Atashi, M. Mansouri, Journal of Industrial and Engineering Chemistry 18 (2012) 1515. [22] B.H. Davis, Fuel Processing Technology 71 (2001) 157. [23] H. Chen, A.A. Adesina, Applied Catalysis A 112 (1994) 87. [24] J. Chang, L. Bai, B.T. Teng, R. Zhang, J. Yang, Y.Y. Xu, H.W. X iang, Y.W. Li, Chemical Engineering Science 62 (2007) 4983. [25] A.Y. Khodakov, W. Chu, P. Fongarland, Chemical Reviews 107 (2007) 1692.