J. Phys. Chem. C 2011, 115, 961–971
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Reaction Kinetics of Ethylene Glycol Reforming over Platinum in the Vapor versus Aqueous Phases† Shampa Kandoi,‡,| Jeff Greeley,‡,§ Dante Simonetti,‡,⊥ John Shabaker,‡,# James A. Dumesic,*,‡ and Manos Mavrikakis*,‡ Department of Chemical and Biological Engineering, UniVersity of Wisconsin-Madison, Madison, Wisconsin 53705 and Center for Nanoscale Materials, Argonne National Laboratory, 9700 South Cass AVenue, Argonne, Illinois 60439 ReceiVed: May 6, 2010; ReVised Manuscript ReceiVed: July 7, 2010
First-principles, periodic, density functional theory (DFT) calculations are carried out on Pt(111) to investigate the structure and energetics of dehydrogenated ethylene glycol species and transition states for the cleavage of C-H/O-H and C-C bonds. Additionally, reaction kinetics studies are carried out for the vapor phase reforming of ethylene glycol (C2H6O2) over Pt/Al2O3 at various temperatures, pressures, and feed concentrations. These results are compared to data for aqueous phase reforming of ethylene glycol on this Pt catalyst, as reported in a previous publication (Shabaker, J. W.; et al. J. Catal. 2003, 215, 344). Microkinetic models were developed to describe the reaction kinetics data obtained for both the vapor-phase and aqueous-phase reforming processes. The results suggest that C-C bond scission in ethylene glycol occurs at an intermediate value of x (3 or 4) in C2HxO2. It is also found that similar values of kinetic parameters can be used to describe the vapor and aqueous phase reforming data, suggesting that the vapor phase chemistry of this reaction over platinum is similar to that in the aqueous phase over platinum. Introduction Diminishing petrochemical resources are increasingly resulting in disruptions in fuel supply and increases in price, adding urgency to the search for sustainable and environmentally safe energy sources. In this respect, biomass has started to play a role of increasing importance because it is renewable, and technologies based on biomass are thus sustainable. Additionally, biomass processing does not contribute to increases in the amounts of CO2 in the atmosphere. Moreover, the annual sustainable growth of biomass worldwide is estimated to be equivalent, on an energy basis, to approximately 50% of the petroleum used each year by the world economy.1 We have recently reported that aqueous phase reforming (APR) of biomass-derived oxygenated hydrocarbons having a C/O stoichiometry of 1:1 (e.g., methanol, ethylene glycol (EG), glycerol, sorbitol, and glucose) over supported Pt catalysts can produce H2 at relatively low temperatures.2–12 This process uses renewable feedstocks and eliminates the need to vaporize the feed, thereby decreasing the energy requirements of the process. Furthermore, the production of CO2 and H2 in APR proceeds via a single catalytic step and produces low levels of CO, a poison to the anode of low temperature fuel cells.13 Reforming of oxygenated hydrocarbons such as EG takes place according to the following stoichiometric reaction: †
Part of the “Alfons Baiker Festschrift”. * To whom correspondence should be addressed. E-mail: (J.A.D.)
[email protected]; (M.M.)
[email protected]. Phone/Fax: 608262-9053. ‡ University of Wisconsin-Madison. § Argonne National Laboratory. | Present Address: UTC Power Corporation, South Windsor, CT 06074. ⊥ Present Address: Department of Chemical Engineering, University of California-Berkeley. # Present Address: BP, Naperville, IL 60563.
C2H6O2 + 2H2O f 2CO2 + 5H2. This reaction is believed to occur via the formation of CO which is then converted to CO2 via the water gas shift (WGS) reaction (CO + H2O f CO2 + H2).14–17 This reaction pathway implies cleavage of C-C and C-H/O-H bonds with subsequent conversion of CO by WGS. Some of the undesired reactions proceed through the formation of alkanes via either parallel pathways involving C-O bond scission or series pathways involving hydrogenation of adsorbed CO and CO2.3,4,6,7 Elementary reactivity studies (e.g., involving in situ spectroscopic measurements) are generally challenging to carry out in condensed phases, and it is thus of interest to understand how experiments and modeling of vapor phase heterogeneous processes, where such elementary studies are feasible, compare to the corresponding aqueous processes. In the present paper, therefore, we compare and contrast the vapor phase and aqueous phase reforming of EG over platinum. We report experimental data for the vapor phase reforming of EG over a Pt/alumina catalyst, where we carried out reaction kinetics measurements at different temperatures, pressures, and feed concentrations. In addition, we report results from first-principles density functional theory (DFT) calculations to determine the potential energy diagram for the interaction of EG with a model Pt(111) surface. We then use this information to construct a reaction scheme for EG reforming over Pt in the vapor phase, and we use this approach to build a microkinetic model that describes the vapor phase reaction kinetics data. We subsequently compare results from our scheme to reaction kinetics data published earlier for aqueous phase reforming of EG.6,7 These models of the vapor phase and the aqueous phase EG reforming reactions on Pt fit the experimental data and provide insight into the similarities and differences in the reforming chemistries occurring in these two phases.
10.1021/jp104136s 2011 American Chemical Society Published on Web 08/12/2010
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TABLE 1: Surface Intermediates Resulting from C-H/O-H and C-C Bond Scission in EG on Pt(111)a binding configuration
species
Egas-phase (eV)
BE (eV)
HOCH2CH2OH
-1281.34
HOCH2CHOH HOCH2CH2O
-1263.69 -1263.31
HOCH2COH HOCHCHOH HOCH2CHO HOCHCH2O OCH2CH2O
-1246.41 -1248.4 -1248.65 -1246.23 -1245.5
HOCHCOH HOCH2CO HOCCH2O HOCHCHO OCHCH2O
-1229.96 -1230.98 -1228.58 -1231.76 -1230.45
HOCCOH HOCHCO HOCCHO OCCH2O OCHCHO
-1213.6 -1214.88 -1213.32 -1211.08 -1215.38
HOCCO OCCHO
-1197.3 -1198.24
OCCO
-1177.88
CH3O CH2O CHO CO CH2OH CHOH COH
-638.63 -623.66 -606.08 -591.15 -638.96 -621.37 -604.31
C2H6O2 species -0.57 monodentate - bound through oxygen C2H5O2 species -1.99 bidentate - bound through carbon and oxygen -1.59 bidentate - bound through two oxygen atoms C2H4O2 species -3.04 monodentate - bound through carbon -1.06 bidentate - bound through two carbon atoms -2.58 bidentate - bound through carbon and oxygen -2.28 bidentate - bound through two oxygen atoms C2H3O2 species -3.44 bidentate - bound through two carbon atoms -2.22 bidentate - bound through carbon and oxygen -4.12 bidentate - bound through carbon and oxygen -1.53 bidentate - bound through two oxygen atoms C2H2O2 species -3.51 bidentate - bound through two carbon atoms -2.06 bidentate - bound through two carbon atoms -3.24 bidentate - bound through carbon and oxygen -3.73 bidentate - bound through two oxygen atoms -0.35 bidentate - bound through two oxygen atoms C2HO2 species -3.62 bidentate - bound through two carbon atoms C2O2 species -6.58 bidentate - bound through two carbon atoms C1 species -1.48 monodentate - bound through oxygen -0.39 bidentate - bound through carbon and oxygen -2.29 monodentate - bound through carbon -1.90 monodentate - bound through carbon -1.88 monodentate - bound through carbon -3.20 monodentate - bound through carbon -4.53 monodentate - bound through carbon
H
-13.62
-2.72
n/a
a Each set of intermediates (delineated by inset labels) corresponds to isomers of a given stoichiometry. Egas-phase refers to the energy of the species in the gas phase. Binding Energy (BE) ) Eslab+adsorbate - Eclean slab - Eadsorbate(g) and corresponds to 1/9 ML coverage. 1 eV ∼ 96.5 kJ mol-1.
Methods Density Functional Theory (DFT) Calculations. DFT calculations are performed using DACAPO, a total energy calculation code.18,19 A two-layer slab of Pt, periodically repeated in a super cell geometry, with five equivalent layers of vacuum between any two successive metal slabs, is used. A 3 × 3 unit cell (surface coverage of 1/9 ML) is employed. Calculations are performed using the PW91 functional; other computational details can be found in a previous publication.20 The metal atoms are kept fixed in their bulk-terminated positions, and all adsorbate atoms are allowed to relax. Convergence with respect to the number of metal layers was confirmed to within 0.1 eV on a three-layer static Pt slab. The lattice constant for bulk Pt is found to be 4.00 Å, in good agreement with the experimental value of 3.91 Å.21 Microkinetic Model. Formulation. A simplified microkinetic model, inspired by results from the DFT calculations, was developed to understand the vapor and aqueous phase reforming kinetics of EG on platinum. A complete microkinetic model would be intractable because, as shown in Tables 1 and 2, the decomposition pathway of EG could consist of numerous intermediates, together with 63 elementary reactions (21 C-C scission and 42 C-H/O-H scission steps) and nearly 20 water gas shift reaction steps.14 Hence, we lump various elementary steps, yielding a 7-step reaction mechanism, as shown below
C2H6O2 + 2* T C2HxO2** +
(6 - x) H2 2
(1)
x C2HxO2** f 2CO* + H2 2
(2)
1 H2O + * T OH* + H2 2
(3)
1 CO* + OH* T CO2 + H2 + 2* 2
(4)
C2H6O2 + * T C2HyO2* +
(6 - y) H2 2
(5)
H2 + 2* T 2H*
(6)
CO + * T CO*
(7)
A* and B** denote monodentate and bidentate adsorbed species. Step 1 is the dehydrogenation of EG to yield a surface intermediate C2HxO2** (x will be defined below), which
Reaction Kinetics of C2H6O2 Reforming over Pt
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TABLE 2: Energetics (in eV) with Respect to Gas-Phase EG and Clean Slab for C-H/O-H and C-C (included in italics) Bond Breaking Steps in EG on Pt(111)a surface reaction HOCH2CH2OH
HOCH2CHOH + H
f HOCH2CHOH + H
initial state -0.57
0.10
1.00
-0.35
f HOCH2COH + 2H
-0.68
0.52
-0.13 -0.41
f HOCH2CHO + 2H
0.10
1.29
0.89 0.11
-0.13
-1.07
f HOCH2CO + 3H
0.06
-0.87
f HOCCH2O + 3H
0.55
-0.37
f CH2OH + COH + 2H
0.53
-1.02
-0.33 0.23 0.86 -0.17
-1.07 -0.55 -0.47 -0.87
0.12 0.80
-0.55 0.35
0.92
-0.54
f HOCHCOH + 3H
f HOCCH2O + 3H
HOCHCOH + 3H
f HOCCOH + 4H f HOCHCO + 4H f HOCCHO + 4H f CHOH + COH + 3H f HOCHCO + 4H
f 2CH2O + 2H
-0.79
-0.79
-0.30
-0.13
f CH2OH + CO + 3H f HOCCHO + 4H
HOCCO + 5H OCCHO + 5H OCCO + 6H a
-0.55 0.35
1.43
0.06
1.25
0.35
2.11
0.58
-1.07
-0.15 0.03 0.19 0.48 -0.11
-1.11 -0.95 -0.56 -1.07 -0.95
-0.87
-0.37
1.19
1.19
-0.03
-1.56
0.48
-0.56
1.19
1.19
f COH + CH2O + 3H
1.03
-0.55
-0.19 0.16 0.45 0.89 1.69
-0.95 -0.56 0.27 -0.60 1.19
1.04 1.43
0.27 -0.08
-0.28 -0.11 -0.41 0.17 -0.09 -0.25 0.23 0.36 1.19
-1.26 -1.68 -1.26 -0.72 -1.62 -1.26 -0.72 -1.21 -0.72
1.19
-1.10
0.27 0.76 0.30 -0.64 -0.27 -0.21 -1.08
-0.72 -0.73 -1.13 -2.23 -1.13 -1.75 -2.77
f HOCHCO + 4H f HOCCHO + 4H f OCHCHO + 4H f CHOH + CHO + 3H f OCCH2O + 4H
f HOCCO + 5H f 2COH + 4H f HOCCO + 5H f OCCHO + 5H f CHOH + CO + 4H f HOCCO + 5H f OCCHO + 5H f COH + CHO + 4H f OCCHO + 5H
-0.55
0.35
-1.11 -0.95
-0.56
1.19
f CH2O + CO + 4H OCHCHO + 4H
-0.37
0.36 0.95
f OCCH2O + 4H
f OCHCHO + 4H f CH2O + CHO + 3H
OCCH2O + 4H
0.52
0.89
f OCCH2O + 4H
HOCCHO + 4H
-0.13
1.62
f CHOH + CH2O + 2H
HOCHCO + 4H
-0.30
0.66
f CH2OH + CH2O + H
f OCHCH2O + 3H
HOCCOH + 4H
0.51
f OCH2CH2O + 2H
OCH2CH2O + 2H
OCHCH2O + 3H
-0.79 -0.30
0.93
f HOCHCHO + 3H f OCHCH2O + 3H
HOCHCHO + 3H
-0.79
0.03 0.27
f CH2OH + CHOH + H
f CH2OH + CHO + 2H
HOCCH2O + 3H
0.03
f HOCHCH2O + 2H
f HOCHCHO + 3H f OCHCH2O + 3H
HOCH2CO + 3H
-0.68
f 2CH2OH
HOCHCHOH + 2H f HOCHCOH + 3H f HOCHCHO + 3H f 2CHOH + 2H HOCH2CHO + 2H f HOCH2CO + 3H
HOCHCH2O + 2H
final state
0.54
f HOCHCH2O + 2H
HOCH2COH + 2H
0.00
f HOCH2CH2O + H
f HOCHCHOH + 2H f HOCH2CHO + 2H
HOCH2CH2O + H
transition state
f OCCHO + 5H f 2CHO + 4H f OCCO + 6H f COH + CO + 5H f OCCO + 6H f CHO + CO + 5H f 2CO + 6H
0.27 -1.26 -0.72 -1.13
Products of bond breaking events adsorb on separate slabs. Initial state entries in bold show the energy of the most stable surface intermediate for each C2HxO2 (x ) 0 to 6) isomeric set. Transition state energies are calculated using the correlations described in the text. Transition state entries in bold show the most stable C-H/O-H and the most stable C-C transition state within each isomeric set. 1 eV ∼ 96.5 kJ mol-1.
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subsequently decomposes into CO and H2 in Step 2. Steps 3 and 4 are the WGS reaction steps, where CO reacts with H2O to yield CO2 and H2. In Step 5, EG decomposes to yield a spectator species (C2HyO2*). Finally, Steps 6 and 7 are the adsorption-desorption of H2 and CO, respectively. We note here that we can write steps 1-5 in terms of gaseous H2 or in terms of H-atoms, because H2 and H* are in quasi-equilibrium (Step 6). We have chosen to write these steps in terms of gaseous H2 because the kinetic parameters for these steps are then independent of the binding energy of H-atoms on the Pt surface. Also, we have written Step 1 as proceeding directly to species C2HxO2** without showing species with intermediate values of x. This lumping is appropriate because we assume that all of these intermediate species are in quasi-equilibrium, and C2HxO2** is the species with the lowest barrier for breaking the C-C bond. Parameterization. Thermodynamic Parameters. The equilibrium constant, Ki,eq ) exp((∆Soi /R) -(∆Hoi /RT)) for each step (except 2, which is assumed irreversible) is calculated by estimating the standard entropy (∆Soi ) and enthalpy change (∆Hoi ) from the thermodynamic properties of the gaseous and o ) for various surface species. The gas phase entropy (Sgas species is taken from standard handbooks, and the entropy of the surface species is calculated using Ssurface ) FlocSloc ) Floc(Sogas - Sotrans,3D), as described in detail elsewhere.22 Values of ∆Hoi for steps 1, 3, 5, 6, and 7 were taken as parameters in the model. ∆H4 was calculated using ∆H4 ) ∆HWGS ∆H3 - ∆H7 where ∆HWGS ) -40 kJ mol-1. According to the above strategies, the model is based on the overall thermodynamic properties of the reactants and products (i.e., ethylene glycol, water, CO2, and H2), thus ensuring that the model achieves thermodynamic consistency for the overall reaction. In addition, we use physically realistic values of binding energies and entropies for adsorbed species. Kinetic Parameters. We parametrize the kinetic model in terms of forward rate constants and then use the equilibrium constants to calculate the reverse rate constants. For the adsorption-desorption Step 7, the adsorption rate constant is defined using collision theory, as described elsewhere.22 For the surface reactions (1 through 4), the rate constants are written as ki ) A exp((-Ei/RT)), where A is the preexponential factor, and Ei is the activation energy barrier (taken as a parameter in the model). For Steps 1 and 3, A is taken as 107 atm-1 s-1, whereas for surface Steps 2 and 4, A is taken as 1013 s-1. Steps 5 and 6 are considered quasiequilibrated, and hence the specific value of Ei is of no kinetic significance for those steps. Having calculated the forward rate constants, we use the equilibrium constants for the respective steps (Ki,eq) to calculate the corresponding reverse rate constants. Summary of Vapor Phase Model DeWelopment. The reaction scheme outlined previously leads to a kinetic model with ten equations and unknowns. We have assumed that the reactor operates as a CSTR. Therefore, we solve five steady-state equations for the gaseous molar flow rates (F) of reactants and products (EG, H2O, CO, CO2, H2), together with five equations for the fractional surface coverages (θ) of adsorbed species (C2HxO2, OH, CO, H, and C2HyO2):
Kandoi et al.
FEG ) FEGin - r1 - r5 FH2O ) FH2Oin - r3 FCO ) FCOin - r7 FCO2 ) FCO2in + r4 FH2 ) FH2in +
r3 r4 (6 - x) x r1 + r2 + + + 2 2 2 2 (6 - y) r5 - r6 2 dθC2HxO2 dt
) r1 - r2
dθOH ) r3 - r4 dt dθCO ) 2r2 - r4 + r7 dt dθH ) 2r6 dt dθC2HyO2 dt
) r5
where F is defined as mol/mol of catalyst/sec, and ri is the net rate of step “i”. Also, we use an algebraic equation for the site balance (i.e., the sum of all surface coverages plus vacant sites is equal to unity). The microkinetic model is used to fit the vapor phase kinetic data obtained under varying conditions of temperature, pressure, and feed composition, as discussed in the experimental section. Summary of Aqueous Phase Model DeWelopment. The same reaction scheme, thermodynamic parameters, and kinetic parameters are used for the aqueous phase model. Shabaker et al. note that gas bubbles form within the liquid-phase flow reactor during APR of EG and that the total pressure inside these bubbles is approximately equal to the total pressure of the system; the bubbles contain water and EG at their respective vapor pressures in the feed solution.6 This formation of gas bubbles within the liquid complicates the aqueous phase model by introducing two additional unknowns as compared to the vapor phase model, leading to a kinetic model with twelve unknowns. Similar to the vapor phase model, we have assumed that the reactor operates as a CSTR, and we solve the above 10 steady-state equations for the gaseous flow rates and coverages. The two additional unknowns (i.e., the flow rate of liquid water out of the reactor (FH2Oliq) and the composition of EG in the liquid effluent (xliq)) are determined based on the vapor pressures of H2O (VPH2O) and EG (VPEG) in the feed at the reaction temperature.
Reaction Kinetics of C2H6O2 Reforming over Pt
xliq + F (1 - xliq) H2Oliq xH2 VPEG (FH2O - FH2Oliq) 1 - xH2 VPH2O
FEG ) FEGliq + FEGgas )
( )( )
Ptot ) pCO + pH2 + pCO2 + pHe + pEG + pH2O )
Ptot (F + FH2 + FCO2 + FHe + xliq*VPEG + Fgas CO (1 - xliq)*VPH2O
where
Fgas ) FCO + FH2 + FCO2 + FHe + FEGgas + FH2Ogas and
FH2Ogas ) FH2O - FH2Oliq We used the microkinetic model to fit experimentally measured aqueous phase kinetic data collected by Shabaker et al.6 Experimental Section The 3 wt % Pt/Al2O3 catalyst used in this study was prepared via incipient wetness impregnation of Catapal-B Al2O3 (Condea) with an aqueous-solution of tetra-amine platinum(II) nitrate (Strem Chemicals). The catalyst was dried overnight at 373 K and calcined in 20% O2 in He at 533 K for 3 h. Prior to reaction kinetics testing and gas adsorption measurements (i.e., CO chemisorption), the Pt/Al2O3 catalyst was reduced in flowing H2 (180 cm3(NTP) min-1) for 2 h (ramped at 0.5 K min-1). The number of active catalytic sites was taken to be equivalent to the irreversible CO uptake at 300 K measured on a gas adsorption apparatus.23 The methods for collecting reaction kinetics data for aqueous phase reforming (APR) of EG are described elsewhere.6 In the present study, reaction kinetics data for vapor phase reforming of EG were collected using a down-flow apparatus described in a previous study.24 The dried, unsieved catalyst (0.209 g) was loaded into a 1/4 in., stainless steel, tubular reactor. Fused SiO2 granules (1.16 g; -4 +16 mesh; Sigma-Aldrich) were loaded upstream of the catalyst bed to aid in vaporization of the feed solution. Type-K thermocouples (Omega) were attached to the outside of the reactor to measure reactor temperature, which was controlled with a series 16 type temperature controller (Dwyer Instruments). Flow rates of H2 (reduction), N2 (carrier gas), H2/N2, and CO/N2 mixtures (H2 and CO reaction order studies) were controlled via mass-flow meters (5850 Brooks Instruments). An HPLC pump (Model 301, Alltech) was used to introduce water and EG feed solutions into the liquid injection/vaporization unit above the reactor.24 The effluent from the reactor was water-cooled in a double-pipe heat exchanger and drained periodically for GC analysis (Agilent 6890 with an FID detector and HP-Innowax column). The effluent gas was analyzed via online gas chromatography (Carle GC (series 8700) with a TCD and HP-5890 GC with TCD and Haysep D 100/ 120 column) and with a Siemens Ultramat 5E CO detector. Feed
J. Phys. Chem. C, Vol. 115, No. 4, 2011 965 solutions were made with EG (spectrophotometric grade; SigmaAldrich) and deionized water. Studies of the effect of temperature on the reforming of EG were conducted at atmospheric pressure with a liquid flow rate of 0.08 cm3 min-1 of 5 wt %, 63 wt %, and pure EG. During these studies, N2 carrier gas was co-fed to the reactor at 230 cm3(NTP) min-1. Similarly, the effect of total pressure was studied at 483 K with 0.08 cm3 min-1 of each of the three feed concentrations and using 230 cm3(NTP) min-1 of N2 carrier gas. To determine reaction order, the liquid feed flow rate, feed concentration, and the flow rates of carrier gas or co-feed gas were adjusted to maintain a constant total molar inlet flow rate of 12 × 103 µmol min-1. The ethylene glycol conversion was maintained below 3% to achieve differential reaction conditions. All reaction order studies were conducted at atmospheric pressure and 483 K. Before data collection at each experimental condition, the catalyst was treated in flowing H2 at 180 cm3(NTP) min-1 at the reaction temperature. We used the Madon-Boudart test in an earlier publication6 to show that the rate of ethylene glycol reforming over supported platinum catalysts was kinetically controlled under the experimental reaction conditions used in the present study. Therefore, effects of transport limitations are not important under our experimental conditions. We have also reported the product distribution of ethylene glycol reforming in our previous paper.6 The only species present in the liquid phase are small amounts of methanol and ethanol (about 150 ppm of each), and smaller amounts of glycolaldehyde and glycolic acid (about 10 ppm each). Ethylene glycol reforming shows essentially 100% selectivity to H2 and CO2 with the amount of CO being about 300 ppm. We have used the Pt dispersion to calculate rates of ethylene glycol reforming per surface site (i.e., turnover frequency). In addition, we showed in our previous publication6 that ethylene glycol reforming is not structure-sensitive over supported Pt catalysts. Results and Discussion Structure and Adsorption Thermochemistry of Reaction Intermediates. Different sites were explored to find the most stable binding configuration for each adsorbate resulting from the dehydrogenation of EG on Pt(111). The energetically most stable adsorption geometry of each surface intermediate is shown in Figure 1. The gas phase energies, binding energies (with respect to clean slabs and species in the gas phase), and binding configurations are given in Table 1; the more negative the binding energy, the stronger the interaction of the adsorbate with the surface. Table 2 shows an alternative presentation of these energetics; the energy of each intermediate with respect to EG in the gas phase and the clean Pt(111) slab is reported in the IS column, permitting a direct comparison of the stability of various intermediates within each isomeric set. In that table, the H-atoms removed from the various reaction intermediates are located on separate Pt(111) slabs. EG (HOCH2CH2OH) binds to Pt(111) with a BE of -0.57 eV, which is slightly stronger than the binding energies of methanol and ethanol on Pt(111).20,25–27 As shown in Figure 1a, the molecule binds through the oxygen atom and stays fairly far from the surface. Dehydrogenation of EG leads to the formation of two possible C2H5O2 species, HOCH2CHOH, and HOCH2CH2O. The adsorbed di-OH species is more stable than the OH-O species (Table 1), and their geometries are shown in Figure 1b. Figure 1c shows the geometries of five possible C2H4O2 surface intermediates that can be formed by the removal of an H atom from the above two C2H5O2
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Figure 1. Cross section and top view of the most stable binding configuration for intermediates resulting from the C-H/O-H and C-C bond scission in EG on Pt(111). Stable adsorbed intermediates for (a) C2H6O2 species, (b) two isomeric C2H5O2 species, (c) five C2H4O2 species, (d) five C2H3O2 species, (e) five C2H2O2 species, (f) two C2HO2 species, (g) C2O2, and (h) various C1 intermediates.
species. They are HOCH2COH, HOCHCHOH, HOCH2CHO, HOCHCH2O, and OCH2CH2O. The two adsorbed di-OH species are the most stable, followed by the OH-O and di-O species. Removal of an H atom from C2H4O2 species results in the formation of five possible C2H3O2 surface intermediates. These
intermediates are shown in Figure 1d. Again, the adsorbed di-O intermediate (HOCHCOH) is the most stable, followed by HOCH2CO, HOCHCHO, and HOCCH2O (the OH-O species). The di-O intermediate (OCHCH2O) is the least stable. As listed in Table 1, HOCHCHO is a tridentate species whereas all the other C2H3O2 species are bidentate.
Reaction Kinetics of C2H6O2 Reforming over Pt
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Figure 2. Gibbs free energy diagram for EG decomposition on Pt(111) relative to EG in the gas phase and a clean Pt(111) slab at 483 K and standard pressures. The energetics of hydrogen resulting from dehydrogenation are referenced to H2(g). Blue diamonds show the most stable intermediates, green circles show the most stable C-H/O-H bond scission transition states, and red inverted triangles show the most stable C-C bond scission transition states within each isomeric C2HxO2 (x ) 0-6) set, as given in Table 2. Note that the most stable transition state does not necessarily correspond to the most stable intermediate. The location of C-H bond breaking is denoted by a small arrow below/above the appropriate hydrogen atom.
Another five surface intermediates can be formed by the subsequent dehydrogenation of the above C2H3O2 species. As shown in Figure 1e, all of these C2H2O2 surface intermediates are bidentate, and the decreasing order of their relative stability on Pt(111) is HOCCOH, HOCHCO, HOCCHO, OCHCHO, and OCCH2O. Removal of an H atom from C2H2O2 species yields two possible C2HO2 species; HOCCO and OCCHO. The adsorbed HOCCO species is more stable and is bidentate, whereas the adsorbed OCCHO is a tridentate species (Figure 1f). Figure 1g shows the OCCO surface intermediate, which is formed as a result of the removal of the last H atom from the C2HO2 species. This OCCO species is bidentate and binds fairly strongly to Pt(111) (see Table 1). Any of the above C2HxO2 (x ) 0 to 6) surface intermediates can undergo C-C bond scission. The resulting C1 surface species (i.e., CH3O, CH2OH, CH2O, CHOH, CHO, COH, and CO) are shown in Figure 1h. Previous DFT literature describing methanol decomposition on Pt(111) has discussed these species in detail, since they are also formed during the dehydrogenation of methanol (CH3OH).20,27 The small difference between the BE (∼0.1 eV or less) reported in previous studies and what we report here is due to the difference in the number of metal layers used in the slabs and lies within the error bars of our DFT calculations. Correlations for Estimating the Activation Energy Barriers. We used Brønsted-Evans-Polanyi type correlations,28 developed for C-H/O-H and C-C bond scissions in other oxygenated compounds on Pt(111),25,27 to estimate the activation barriers for bond cleavage reactions in EG on Pt(111). These correlations relate the energy of the transition state to the energy of the products (final state) with each surface reaction being defined in the exothermic direction. The energies of the transition states (TS) and final states (FS) are relative to the energy of the corresponding initial state in the gas-phase. Greeley and Mavrikakis studied methanol decomposition on Pt(111) using DFT calculations and found that the energies of the transition states for elementary steps in methanol decom-
position could be related to the respective final states energies through the equation:27 ETS (eV) ) 0.90EFS (eV) + 0.61. We use this correlation to estimate the transition states for the C-H/ O-H bond scission steps in EG on Pt(111). Similarly, Alcala et al. studied ethanol decomposition on Pt(111). Their results25 show a linear correlation between the energies of the transition states for C-C and C-O bond scission elementary steps and the energies of the respective final states: ETS(eV) ) 0.97EFS(eV) + 1.45. We use this correlation to estimate the barriers for C-C bond breaking in EG on Pt(111). The slopes of these correlations suggest that the transition states for these elementary steps resemble the final states (written in the exothermic direction). Estimated Potential Energy Surface (PES) for Ethylene Glycol Decomposition. Table 2 shows the energetics of the initial state (IS) and final state (FS) for all possible C-H/O-H and C-C bond breaking steps in EG on Pt(111). The energies of the transition states for these bond scissions are calculated using the above linear correlations. The Gibb’s free energy changes (at 483 K) for these reaction steps (relative to EG in the gas phase and the clean slab) were also calculated and are reported in Figure 2. To determine the enthalpy change for each step, DFT results reported in Table 2 were used, and thermal and zero point energy corrections were assumed to be small. To estimate the standard entropy change for each step, we assumed that (i) dehydrogenation leads to the loss of gas-phase translational entropy for EG (calculated using the standard translational entropy of a gaseous molecule with three degrees of freedom) and (ii) an adsorbed transition state has the local entropy (i.e., rotation and vibration) of the corresponding C2HxO2 species. In these calculations, we allow the adsorbed H-atoms formed during dehydrogenation to desorb as H2 into the gas phase (because the hydrogen adsorption/desorption step is assumed to be quasi-equilibrated). The enthalpy and entropy of H2 in the gas phase are taken as -31.8 eV (from DFT) and approximately 131 J mol-1 K-1 (from standard handbooks), respectively. The most stable species within each isomeric set and the most stable transition states for C-H/O-H and C-C
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Figure 3. Comparison of model predictions (b) and experimental data (9) for the effects of a) total system pressure and b) EG feed concentration on H2 production during aqueous phase reforming of EG. Experimental data from Shabaker, et al.6
bond scission are included in Figure 2. The DFT calculations suggest that the C-C bond in EG is broken at an intermediate value of x, that is, around C2H3O2; at this value, the barriers to cleave C-C bonds become approximately equal to those for C-H/O-H cleavage, implying that C-C cleavage becomes a kinetically competitive pathway at this point. Experimental Results The nominal loading for the catalyst used in this study was 3.0 wt % Pt. Previous work documented the effects of total pressure and feed concentration on H2 production during the aqueous-phase reforming of EG,6 and these data are shown in Figure 3. Figure 3, as well as subsequent figures, also present values of turnover-frequencies predicted by our kinetic model, and these results will be described later in this paper. Shabaker et al. found that the total pressure (Figure 3a) has a strong inhibiting effect on the rate of H2 production, and that this effect is more pronounced as the pressure of the system is decreased closer to the bubble point of the feed solution.6 For the purposes of the study in this paper, the order with respect to total pressure is taken to be -1.7 (as shown in Figure 3a). The reaction order with respect to EG (Figure 3b) is 0.45 over the entire range of EG concentration.6 Figure 4 shows the effect of temperature on the rates of H2 and CO2 production for the vapor phase reforming of various concentrations of water-EG feed solutions. The slopes of the semilog plots in Figure 4a correspond to activation barriers of 70 and 110 kJ mol-1 for H2 and CO2 formation, respectively, for a 5 wt % EG feed solution. When more concentrated EG solutions were fed to the reactor, smaller amounts of CO2 were detected, and Figure 4b (63 wt % EG) and Figure 4c (pure EG)
Figure 4. Comparison of model predictions and experimental data for the effect of varying temperature on H2 and CO2 production during vapor phase reforming of EG. Experimental H2 TOF (9) and CO2 TOF (2), compared with model predictions of H2 TOF (b) and CO2 TOF ((), at various temperatures for EG feed concentrations of (a) 5, (b) 63, and (c) 100 wt % at 1 bar and liquid feed flow rates of 0.08 cm3 min-1.
thus show only the values of H2 TOF and corresponding activation barriers of 60 and 80 kJ mol-1, respectively. Increasing the feed concentration has little effect on the formation of H2. However, the CO2 turnover frequencies are an order of magnitude lower at the higher EG feed concentration (63 wt % compared to 5 wt %). This decrease in the rate of CO2 production results from less water to react with CO via WGS. Figure 5 shows the results obtained by varying the total system pressure. Increasing the total pressure slightly inhibits the rate of H2 formation, as evidenced by the negative reaction orders for 5 wt %, 63 wt %, and pure EG feeds. However, for 5 and 63 wt % solutions, the reaction order is fractional (0.5) for the rate of CO2 formation. Also, the values of TOF for CO2 formation are almost an order of magnitude lower for the 63 wt % feed solution compared to the 5 wt % feed because of
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Figure 5. Comparison of model predictions and experimental data for the effect of varying total system pressure on H2 and CO2 production during vapor phase reforming of EG. Experimental H2 TOF (9) and CO2 TOF (2), compared with model predictions of H2 TOF (b) and CO2 TOF ((), at various total pressures for EG feed concentrations of (a) 5, (b) 63, and (c) 100 wt % at 483 K and liquid feed flow rates of 0.08 cm3 min-1.
Figure 6. Effect of reactant partial pressure during vapor phase reforming of EG. Experimental H2 TOF (9) and CO2 TOF (2) compared with model predictions of H2 TOF (b) and CO2 TOF (() at various partial pressures of (a) EG, (b) H2O, and (c) CO at 483 K and 1 bar.
decreased WGS activity at higher EG concentrations. At higher total pressures, the formation of H2 from EG is inhibited; however, the higher total pressures lead to higher partial pressures of H2O, which lead to production of H2 and CO2 from adsorbed CO species via the WGS reaction. Thus, for dilute feed solutions H2 from WGS compensates for H2 lost due to the negative pressure order, resulting in similar H2 turnover frequencies for the two feeds, whereas the rate of CO2 production is higher for the more dilute EG feed because of the higher water partial pressure. The sharp decrease at 10 bar for the rate of H2 production from pure EG could be a result of incomplete vaporization of EG. Figure 6 shows the vapor phase reaction orders with respect to EG (panel 6a), H2O (panel 6b), and CO (panel 6c). The rate of H2 production is fractional order with respect to both EG and H2O. The reaction order with respect to EG (0.4) is the
same as for APR of EG.5 The rate of CO2 production is approximately first order with respect to H2O, indicating that OH does not have a significant site-blocking effect (step 4 in the reaction scheme), and therefore water-derived OH does not bind strongly on surface Pt sites. We note that cofeeding H2 with EG slightly inhibits the rate of H2 production (order of -0.04), whereas cofeeding CO strongly inhibits H2 production. This inhibiting effect of CO is most likely due to increased surface coverage by CO, which blocks adsorption of EG.29 Results of the Microkinetic Models. The values of the kinetic parameters used to describe the reaction kinetics data for both vapor and aqueous phase reforming are summarized in Table 3. We note that since the total number of parameters that needed to be fit is already relatively modest and because some of the values do not correspond to elementary reaction steps, we have not assumed any direct link between the kinetic
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TABLE 3: Fitted Values for Kinetic Parameters and Confidence Limits for the Most Sensitive Parametersa vapor phase kinetic parameter
value
Eact 1 Eact 2 Eact 3 Eact 4 ∆H1 ∆H3 ∆H5 ∆H6 ∆H7
53.7 135.8 80.2 60.0 -40.0 10.0 -55.0 -50.0 -108.7
95% confidence interval (5.64 (0.69 (2.67
(5.83
aqueous phase value 50.0 133.3 75.0 60.0 -39.0 10.0 -55.0 -50.0 -95.0
95% confidence interval (1.35 (3.88
All values are in kJ mol-1. The definition of the various elementary steps associated with the Eact and ∆H is given in the reaction scheme presented in the Microkinetic Modeling section of this paper. a
and thermodynamic parameters via a BEP relation in this case. Each model was found to be sensitive to only a few parameters (highlighted in bold), and the fitted values of those parameters were determined using Athena Visual Workbench engineering software.30 The 95% confidence limits of the fitted parameters are also included in the table. A comparison of the kinetic parameters used to describe the vapor phase and aqueous phase reforming of EG shows that they are similar (difference of less than 15 kJ mol-1). Also, the values of x and y in the simplified reaction mechanism proposed earlier were found to be 4 and 6, respectively, (to yield C2H4O2** and C2H6O2* species) for both the vapor phase and aqueous phase reforming of EG (we note that other values of x and y gave very poor fits to our experimental data, strongly implying that the reported values are physically and chemically relevant). The similarity between the parameters of the vapor and aqueous phase models indicates that, even though the level of complexity involved in the two cases is quite different, similar chemistry occurs on platinum. The value of x suggested from the microkinetic model is in accord with the predictions of DFT calculations that the C-C bond scission in EG on platinum occurs at intermediate values of x. The heats of reaction (∆Hi) for steps 1, 3, 5, 6, and 7 are estimated from DFT calculations on Pt(111) to be about 10, 80, -55, -85, -183 kJ/mol, respectively. The activation energy barrier for step 3 (Eact 3) calculated using DFT is about ∼60 kJ/mol.31 (It is difficult to determine the activation energy barriers of other steps (1, 2, and 4) directly from DFT, because those steps are lumped steps). If we compare these DFT results with the fitted values given in Table 3, we find non-negligible differences in some situations. This discrepancy can be partly attributed to the fact that the above DFT values correspond to a flat (111) surface whereas defect sites might also play an important role in measurements on supported platinum; in some cases, inclusion of surface relaxation effects in the DFT calculations may also reduce the DFT/microkinetic discrepancies. A general idea of the magnitude of such effects can be obtained, for example, by comparing the value of ∆H3 calculated on a stepped Pt(211) surface to that on the flat surfaces (this comparison is made only to point out the general trend, as ∆H3 is not found to be a highly sensitive parameter in the microkinetic model fit); a value of around 24 kJ/mol32 is found, which is closer to the fitted value of this parameter (Table 3). In addition, the high coverage of the most abundant surface intermediate (C2H4O2, as discussed later) might have an influence on the binding energies of other species. For example, significant coverage of C2H4O2 on the surface might destabilize
CO, which otherwise binds to platinum surfaces strongly. This result might justify the relatively weaker binding energy of CO (∆H4) fitted by the model. We note, however, that none of these effects will change our conclusion that C-C bond scission occurs at intermediate values of x. A comparison of the experimental and model predicted values of the TOF for aqueous phase reforming is shown in Figure 3. We find that the aqueous phase model describes the experimental data well. The same comparison for the vapor phase reforming under varied conditions of temperature, pressure, and feed compositions is shown in Figures 4, 5, and 6, respectively. In this case, most of the experimentally observed trends can be captured. However, for a few data points, the model prediction deviates from experimental data, for example, at high temperatures (Figure 4) as well as for the reaction order of EG (Figure 6a). We note that the model predicted the reaction order with respect to H2 (-0.05) well, although the TOF values deviated by a factor of 2, similar to the deviation in Figure 5 at low pressure. The deviation between the model prediction and experimental TOF at the data point corresponding to 10 bar for the pressure study with pure EG (Figure 5c) can be attributed to the model failing to simulate the effect of incomplete vaporization of the feed. In addition to reproducing the experimental data for vapor and aqueous phase reforming of EG on platinum, the model also gives insight into the reforming pathway and the coverage of various species. We find that under most of the vapor phase reforming conditions (except when CO is cofed), the most abundant species is C2H4O2**, followed by smaller coverages of CO* and H*. Similar trends in the relative coverage of various species are observed for aqueous phase reforming although the number of vacant sites is, in general, greater in the case of aqueous compared to vapor phase reforming. The low coverage of CO under reforming conditions is consistent with the fact that CO produced is converted into CO2 via the WGS reaction, and therefore CO does not poison the catalyst surface. Conclusions We have performed periodic DFT calculations to study the decomposition of ethylene glycol (EG) on the Pt(111) surface. The most stable C2HxO2 (x ) 0 to 6) species derived from the subsequent dehydrogenation of EG (with respect to EG in the gas phase and clean slab) are HOCH2CH2OH, HOCH2CHOH, HOCH2COH or HOCHCHOH, HOCHCOH, HOCCOH, HOCCO, and OCCO. Within each isomeric set, the adsorbed di-OH species are found to be most stable, followed by OH-Ocontaining species, and then by di-O species. We have used this information to develop an energy diagram for EG interaction with platinum surfaces, which predicts that C-C bond cleavage occurs for intermediate values of x in C2HxO2. We have also described reaction kinetic data collected for the vapor phase reforming of EG on a supported Pt catalyst, together with the corresponding results for the aqueous phase reforming of EG, which have been reported in previous publications. Finally, microkinetic models have been built to compare and contrast the vapor-phase and aqueous-phase reforming kinetics at the atomic level. The same reaction scheme is used to build the two models, with two additional equations required to build the aqueous phase model to account for the formation of gas bubbles during APR. These models describe the experimental kinetic data reasonably well using similar values of the kinetic parameters, indicating that the vapor phase and aqueous phase reforming chemistry is similar on platinum. The microkinetic
Reaction Kinetics of C2H6O2 Reforming over Pt model also suggests that the C-C bond in EG is broken on platinum at an intermediate value of x in C2HxO2. Acknowledgment. We wish to thank Professor Baiker for his leading work on the kinetics and mechanisms of liquidphase catalytic reactions. This material is based upon work supported as part of the IACT, an Energy Frontier Research Center funded by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences. Calculations were performed in part using supercomputing resources at the following institutions: (1) EMSL, a national scientific user facility located at Pacific Northwest National Laboratory; (2) the National Center for Computational Sciences (NCCS) at Oak Ridge National Laboratory; (3) the Center for Nanoscale Materials (CNM) at Argonne National Laboratory; and (4) the National Energy Research Scientific Computing Center (NERSC). EMSL is sponsored by the U.S. Department of Energy’s Office of Biological and Environmental Research. NCCS, CNM, and NERSC are supported by the Office of Science of the US Department of Energy under Contract No. DE-AC05-00OR22725, DE-AC02-06CH11357, and DE-AC02-05CH11231, respectively. References and Notes (1) Ragauskas, A. J.; Williams, C. K.; Davison, B. H.; Britovsek, G.; Cairney, J.; Eckert, C. A.; Frederick, W. J.; Hallett, J. P.; Leak, D. J.; Liotta, C. L.; Mielenz, J. R.; Murphy, R.; Templer, R.; Tschaplinski, T. Science 2006, 311, 484. (2) Cortright, R. D.; Davda, R. R.; Dumesic, J. A. Nature 2002, 418, 964. (3) Davda, R. R.; Alcala, R.; Shabaker, J.; Huber, G.; Cortright, R. D.; Mavrikakis, M.; Dumesic, J. A. Science and Technology in Catalysis 2002 2003, 145, 79. (4) Davda, R. R.; Shabaker, J. W.; Huber, G. W.; Cortright, R. D.; Dumesic, J. A. Appl. Catal., B 2003, 43, 13. (5) Huber, G. W.; Shabaker, J. W.; Dumesic, J. A. Science 2003, 300, 2075. (6) Shabaker, J. W.; Davda, R. R.; Huber, G. W.; Cortright, R. D.; Dumesic, J. A. J. Catal. 2003, 215, 344. (7) Shabaker, J. W.; Huber, G. W.; Davda, R. R.; Cortright, R. D.; Dumesic, J. A. Catal. Lett. 2003, 88, 1.
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