Combined DFT, Microkinetic, and Experimental Study of Ethanol

Article pubs.acs.org/JPCC

Combined DFT, Microkinetic, and Experimental Study of Ethanol Steam Reforming on Pt Jonathan E. Sutton,† Paraskevi Panagiotopoulou,†,‡ Xenophon E. Verykios,‡ and Dionisios G. Vlachos*,† †

Catalysis Center for Energy Innovation and Center for Catalytic Science and Technology, Department of Chemical and Biomolecular Engineering, University of Delaware, Newark, Delaware 19716, United States ‡ Department of Chemical Engineering, University of Patras, GR-26504 Patras, Greece S Supporting Information *

ABSTRACT: Density functional theory (DFT) calculations for the thermal decomposition and oxidative dehydrogenation of ethanol, mechanistic aspects of water−gas shift reaction, and experimental kinetic data are integrated so as to develop and assess a comprehensive DFT-based microkinetic model of low temperature ethanol steam reforming on Pt catalysts. The DFT calculations show (1) that the C−C scission should occur late in the dehydrogenation sequence, (2) that the C−C scission barriers in highly dehydrogenated intermediates are comparable to early C−H abstraction barriers, and (3) that the oxidative dehydrogenation reactions should not be important under steam reforming conditions. The DFT-parametrized model shows good qualitative agreement with experiments, with reasonable deviations attributed to modeling only the metal chemistry (i.e., excluding support effects). Both the model and the experiments show that the overall mechanism is simply thermal decomposition of ethanol followed by incomplete water−gas shift. The most abundant surface species in the model are the decomposition products CO, H, and free sites, while the key reactive intermediates are present in much lower amounts. Unlike findings of simplified previous models, the rate determining step was identified as the initial dehydrogenation of ethanol, while the selectivity to C1 products is controlled by the C−C cracking of CHCO. Brønsted−Evans−Polanyi (BEP) correlations for the oxidative dehydrogenation reactions are developed and the effect of coadsorption on BEPs is discussed.

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INTRODUCTION Ethanol steam reforming has received a tremendous amount of attention over the past decade because ethanol is readily available, has low toxicity, is compatible with current fuels infrastructure, and has a relatively high H2 content. Indeed, a number of reviews of the topic have recently been published.1−3 It is generally believed (1) that the mechanism depends both on the metal and the support, with the metal contributing to C−C and C−H bond scission and the support participating in the water−gas shift reaction, in dehydration and possibly etherification chemistry, and (2) that the primary deactivation mechanism is either via coke resulting from the decomposition of ethylene on the support (typically γ-Al2O3) or the formation of surface acetates (on reducible oxides, e.g., CeO2). Nevertheless, a number of questions still remain regarding the process. In particular, the exact role of the metal prior to deactivation is still not well understood. The overall reaction of ethanol steam reforming CH3CH 2OH + 3H 2O → 2CO2 + 6H 2

and methane steam reforming (or its reverse methanation reaction) CH4 + H 2O ↔ CO + 3H 2

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It is not entirely clear which mechanism(s) is active in the overall process. Both theoretical (primarily density functional theory (DFT)) and experimental methods have been utilized in an effort to elucidate the mechanism. DFT studies have been carried out for ethanol synthesis and decomposition on a variety of late transition metals, including Ag,4 Au,4 Co,4 Cu,4 Ir,4 Ni,4 Pd,4,5 Pt,4,6,7 Rh,4,8−12 and Ru.7 There are a number of experimental studies of the intrinsic kinetics of ethanol steam reforming but relatively few postulate either a Langmuir− Hinshelwood−Hougen−Watson (LHHW) kinetic model or a power law rate expression. Those that did entailed metals such as Co,13 Ni,14−19 Ni−Pt,20,21 Pt,22 Rh,23−25 and Ru26 on Al2O313,14,16−21,24−27 and CeO222,23,25 based supports (among others). There have also been very few studies to postulate a more detailed kinetic model of ethanol synthesis, decomposition, or steam reforming. The first such model, to our knowledge, was by Ferrin et al.7 who constructed a mean field microkinetic model of ethanol thermal decomposition in the presence of H2 for a number of metals. Subsequent to their work, two separate kinetic Monte Carlo (KMC) simulations of

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consists of at least three submechanisms at low temperatures. These are thermal decomposition CH3CH 2OH → CH3CHO + H 2 → CH4 + CO + H 2 (2)

Received: December 20, 2012 Revised: February 13, 2013 Published: February 13, 2013

water−gas shift CO + H 2O ↔ CO2 + H 2 © 2013 American Chemical Society

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dx.doi.org/10.1021/jp312593u | J. Phys. Chem. C 2013, 117, 4691−4706

The Journal of Physical Chemistry C

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ethanol synthesis were performed.9,12 Most recently, a microkinetic model for the autothermal reforming of ethanol was postulated.28 Recently our group postulated microkinetic models for ethylene glycol (a related oxygenate) decomposition and reforming.29,30 All of the studies to date are lacking in one or more respects. Many of the aspects of the various modeling techniques were recently reviewed.31 DFT studies at the zero coverage limit do not account for the actual rate and the coverage of intermediates of a given elementary reaction. Instead, postulates about the rate determining step are made solely on the basis of which reaction in the mechanism has the highest barrier. While useful, such conclusions can be misleading. Further, most of the kinetic studies proposed LHHW type models with regressed parameters or power law rate expressions. LHHW models with regressed parameters suffer generally from (1) the need to make a priori assumptions about the most abundant surface intermediate (MASI) and the rate determining step (RDS) and (2) the very limited number of reactions (typically no more than a dozen or so) whose parameters can be independently estimated via regression and the ease of algebraically manipulating the equations in order to obtain a closed form rate expression. Mean field microkinetic and KMC models overcome many of the shortcomings of the LHHW approach. Specifically, they eliminate the need for a priori assumptions and the limitations on the number of elementary reactions imposed by the need to generate a closed form rate expression. The published microkinetic and KMC models represent an important first step toward a more thorough kinetic model of ethanol decomposition, synthesis, and reforming, but none is directly applicable to low temperature ethanol steam reforming. There is also some debate as to the importance of various steps in the reforming and decomposition of ethanol. For instance, Ferrin et al.7 assumed that all dehydrogenation reactions were in equilibrium and concluded that the rate controlling step for C−C cleavage occurs in CHCO*. In contrast, more recent microkinetic models of ethylene glycol decomposition and steam reforming29,30 show that the rate determining step is one of the initial H abstractions and that the C−C cleavage reactions were less important for determining the overall rate. In order to achieve a more comprehensive understanding of the low temperature ethanol steam reforming mechanism (including elucidating the relative importance of dehydrogenation and cracking reactions), we propose here the most comprehensive microkinetic model developed to date. Further, we have calculated the vast majority of our parameters with selfconsistent DFT, and by doing that, we eliminate possible issues with model parametrization. In this paper we address a number of questions. First, we explore the conclusions to be drawn from a DFT study of ethanol decomposition and reforming regarding the possible dominant pathways from a purely energetic standpoint. Then we solve our microkinetic model and conduct more detailed analyses in order to gain insight into the key reactions, species, and mechanisms in the lowtemperature ethanol steam reforming with comparisons to our DFT data where relevant. In the process we illustrate why DFT studies alone can lead to misleading conclusions. We also show that the identity of the MASI and the key reactive intermediate are not necessarily the same. We show that, as with ethylene glycol, the rate determining step in ethanol steam reforming is an initial H abstraction. Finally, we assess the model with new experimental data.

Article

METHODOLOGY

DFT Calculations. We employed the SIESTA density functional theory code32 with Troullier−Martins normconserving scalar relativistic pseudopotentials,33 a double-ζ plus polarization (DZP) basis set, and the Perdew−Burke− Ernzerhof GGA functional.34 All slabs were four layers thick, with the two bottom layers frozen in their bulk positions. Metal slabs were separated by 15 Å of vacuum. The Pt lattice constant was taken as 4.02 Å, consistent with prior work in our group using this computational setup.35 Spin polarization was included in the calculations of gas-phase species but neglected otherwise. The majority of the calculations for simple thermal decomposition reactions were carried out on 2 × 2 unit cells of Pt(111), while the calculations for the oxidative dehydrogenation steps were carried out on 3 × 3 unit cells. In the former case, a 5 × 5 × 1 Monkhorst-Pack mesh was used, while in the latter, a 3 × 3 × 1 mesh was employed. Calculations for a few sensitive thermal decomposition reactions (as identified by sensitivity analysis on the microkinetic model later) were repeated on 3 × 3 unit cells with a 5 × 5 × 1 Monkhorst-Pack mesh. Vibrational frequencies were also calculated for these species and transition states in order to obtain temperature corrected thermochemistry and activation energies. This methodological twist (calculations in a smaller cell and refinement in a larger cell) can be thought as an example of hierarchical refinement and saves computational time.31 Transition state searches were carried out using a constrained optimization scheme36 coupled with a local transition state search algorithm.37 Briefly, the length of the bond being formed or broken is constrained at an estimated value and the total energy is minimized with respect to the remaining degrees of freedom. The local transition state search algorithm is capable of updating the constraint on-the-fly as needed to improve convergence to the correct transition state. The transition state is required to meet two additional criteria: (1) all forces on all atoms must vanish and (2) the total energy must be at a maximum along the reaction coordinate but at a minimum in the other degrees of freedom (i.e., at a saddle point in the potential energy surface). This approach has been used successfully in various works.29,35,38,39

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EXPERIMENTAL SECTION Catalyst Preparation and Characterization. The Pt/γAl2O3 catalyst was prepared employing the wet impregnation method using γ-Al2O3 (Alfa products) as support and (NH3)2Pt(NO2)2 (Alfa products) as platinum precursor salt, as reported elsewhere.40 The nominal Pt loading of the catalyst thus prepared was 0.5 wt %. The catalyst was characterized with respect to its specific surface area (BET), exposed metallic surface area, and mean crystallite size employing nitrogen physisorption at liquid nitrogen temperature and selective chemisorption of H2 at 25 °C. Details on the apparatus and procedures used can be found elsewhere.41 Results of catalyst characterization showed that the specific surface area (BET) of the 0.5%Pt/Al2O3 catalyst was 83 m2/g, the Pt dispersion was 99%, and its mean crystallite size (dPt) was 0.9 nm. Catalytic Performance Tests under Steady-State Conditions. Catalytic performance tests were carried out using an apparatus that has been described in detail elsewhere. Briefly, it consists of a flow measuring and control system equipped with four mass-flow controllers (MKS), an HPLC 4692



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pump (Marathon Scientific Systems), and an evaporator for feeding the ethanol/water stream into the quartz microreactor and two gas chromatographs connected in parallel through a common set of switch valves for online analysis of reactants and products. The first GC is equipped with two packed columns (Porapak-Q, Carbosieve) and two detectors (TCD, FID) and operates with He as the carrier gas. The second GC is equipped with a Carbosieve column and a TCD detector. This chromatograph uses N2 as the carrier gas and was solely used for the determination of H2 in the reformate gas. The catalytic performance of the Pt/Al2O3 catalyst for the steam reforming of ethanol has been investigated in the temperature range of 300−400 °C using a feed stream consisting of 12.5% CH3CH2OH, 37.5% H2O, balance He. In a typical experiment, 650 mg of fresh catalyst (0.18 < d < 0.25 mm) is placed in the reactor and reduced in situ at 300 °C for 2 h under a hydrogen flow of 60 cm3/min. After purging with He for 15 min, the sample was exposed to the reaction mixture (120 cm3/min), and ethanol conversion as well as product distribution were determined as a function of reaction temperature. Details on the methods and procedures employed can be found elsewhere.40 The effect of partial pressure of ethanol and steam on the kinetic reaction rate has been investigated at 300 °C using a feed stream consisting of 1.3−12.5% CH3CH2OH and 12.8− 75.3% H2O (balance He). Measurements of intrinsic kinetic rates were obtained under differential reaction conditions, where the conversion of ethanol was typically below 10%, by varying the mass of catalyst (15−34 mg) and the flow rate (530−950 cm3/min). All experiments were performed at near atmospheric pressure. Microkinetic Modeling. The microkinetic model was formulated in two stages. First, a comprehensive list of elementary reactions comprising thermal decomposition was generated: CHxCHyOHz * + * ↔ CHx − 1CHyOHz * + H*

(5)

CHxCHyOHz * + * ↔ CHxCHy − 1OHz * + H*

(6)

CHxCHyOH * + * ↔ CHxCHyO * + H*

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CHxCHyOHz * + * ↔ CHx * + CHyOHz *

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CHxCHyOHz * + * ↔ CHxCHy * + OHz *

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CHxCHy * + * ↔ CHxCHy − 1 * + H*

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CHxCHy * + * ↔ CHx * + CHy *

(11)

CHxOHz * + * ↔ CHx − 1OHz * + H*

(12)

CHxOH * + * ↔ CHxO* + H*

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CHxOHz * + * ↔ CHx * + OHz *

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CHx * + * ↔ CHx − 1 * + H*

(15)

CHxCHyOHz * + OH w * ↔ CHx − 1CHyOHz * + OH w + 1 *

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CHxCHyOHz * + OH w * ↔ CHxCHy − 1OHz * + OH w + 1 *

(17)

CHxCHyOH * + OH w * ↔ CHxCHyO * + OH w + 1 * (18)

were investigated via DFT but were not included in the microkinetic model, as previous work30 has shown these reactions to be practically unimportant on Pt. This assumption was further validated by the current modeling results, as will be discussed later. Excluding oxidative dehydrogenation reactions, there were 160 reactions involving 67 gas and surface species in the microkinetic model. The model was parametrized using DFT-based values from Pt(111) almost exclusively. The modeling of the entire nanoparticle as a single Pt(111) facet is not likely to cause significant errors in the rate because (1) this is probably the dominant facet of the supported nanoparticles in this work and (2) prior KMC work on water−gas shift42 has shown that under similar conditions the steps, edges, and kinks are covered by CO*, likely rendering these sites inactive for thermal decomposition. It would be of interest to quantify this effect via a KMC model, which could explicitly account for the surface heterogeneity present in supported nanoparticles, but such a work is beyond the scope of this paper. The thermochemistry for the needed surface intermediates was taken from previous work.30,43 Pre-exponentials were calculated as (kBT)/h, where kB is Boltzmann’s constant, h is Planck’s constant, and T is the absolute temperature and the rate constant calculated as k=

⎛ −E ⎞ kBT exp⎜ A ⎟ ⎝ RT ⎠ h

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Activation energies were taken from the DFT work performed herein where possible. Most of these activation energies were calculated at 0 K and without the zero point energy correction in order to reduce the computational burden. This does not introduce significant error, as explained in the Supporting Information. We were unable to locate transition states for two reactions, the dehydrogenation of CCH2O* and CCHO*. These activation energies were assumed to be equal to the α C−H dehydrogenations in CCH2OH* and CCHOH*, respectively. This assumption should have little or no effect on the model results owing to the unimportance of these species in the mechanism. Based on preliminary sensitivity analyses of the untuned DFT-based microkinetic model, a number of reactions were identified as potentially sensitive, and vibrational frequency calculations were performed in order to replace the 0 K nonzero point energy corrected values with Gibbs free energies of activation calculated at the actual temperatures of reaction. These reactions were

where a “*” denotes either an adsorbed species (when attached to a molecular fragment) or a vacant surface site (when used alone). Second, a complete water−gas shift mechanism on Pt was taken from the literature.30 Oxidative dehydrogenation type reactions, namely, 4693

CH3CH 2OH* + * ↔ CH3CHOH* + H*

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CH3CH 2O* + * ↔ CH3CHO* + H*

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CH3CH 2OH* + * ↔ CH3CH 2O* + H*

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CH3COH* + * ↔ CH3C* + OH*

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CHCO* + * ↔ CH* + CO*

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CH3CH 2 * + * ↔ CH3CH* + H*

(25)

CH3 * + * ↔ CH 2 * + H*

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focus our attention on trends in binding and activation energies before developing a most energetically favorable pathway for the thermal decomposition reactions. For brevity, graphical representations for structures and transition states of the form C2HxOHy* are given in the Supporting Information. The structures of the smaller species are consistent with those previously published by our group29,35,48 and elsewhere in the literature49 and are not included. Thermal Decomposition Reactions. The binding energies of the stable adsorbates with respect to gas-phase ethanol are given in Table 1. Table 1 also contains the corresponding

All water−gas shift parameters but two were taken directly from the literature30 without alteration. These reactions also utilized Gibbs free energies of activation. The only exceptions were the pre-exponential for the formation of COOH* from CO* and OH* and the pre-exponential for the subsequent decomposition of COOH* to CO2* and H*. These were artificially increased by a factor of 1000 in order to improve agreement between the model and experimental CO x selectivities. These increases are an approximate means of compensating for additional contributions to the water−gas shift rate not included in our model. For example, as previously mentioned, because the rate is dominated by thermal decomposition on the terraces and we fit the surface area to match the conversion at a given temperature, we effectively only account for the surface area of the terraces. This neglects the steps, edges, and kinks, which have been shown via KMC modeling to significantly contribute to water−gas shift rates at similar reaction conditions.42 Additionally, our metal-only model completely neglects the support. Recent DFT calculations by Roy et al.44 on the dehydration of various alcohols on γ-Al2O3 show that (1) water adsorbs very strongly on the support (and more strongly than ethanol), (2) water dissociation to OH* and H* has a relatively low barrier, and (3) OH* is more thermodynamically stable than water on the support. Thus, the effective concentration of OH* available to react with CO* to form COOH* could be much higher than predicted by the current model (e.g., via spillover onto the metal or via a reaction at the metal−support interface). There is probably a complex interplay of effects responsible for the enhanced rate of water−gas shift. In order to elucidate the complete mechanism, it would be necessary to develop models for the chemistry on the metal, on the support, and at the interface. Given the complexity of such a task coupled with the dominance of the metal in the chemistry, as a first approximation, we model primarily the metal and apply these approximate corrections to the sensitive water−gas shift steps. Further work should be done to develop models for the support and the interface. These artificial increases are not enough to affect the rate or conversion of ethanol (as demonstrated later in the sensitivity analysis). Coverage effects for both the stable adsorbates and for the activation energies (following the methodology of Salciccioli et al.45) were incorporated into the model. A complete list of reactions and parameters is provided in the Supporting Information. The model was then solved in an idealized plug flow reactor (PFR) using an in-house reactor code built around the CHEMKIN-II kinetics libraries.46,47 Our reactor code routinely outputs the gas-phase compositions and surface coverages; the forward, reverse, and net rates of each elementary reaction; and the amount a particular reaction contributes to the production or consumption of a species. It can also be used to perform a sensitivity analysis (SA) on the mechanism in order to identify the sensitive reactions. DFT Results. In this section we explore the DFT results for two classes of reactions of ethanol derivatives: thermal decomposition and oxidative dehydrogenation. Due to the large number of species and reactions investigated, we will first

Table 1. Binding Energies (eV) of Stable Adsorbates on Pt(111) with Respect to Gas-Phase Ethanola species

this work

ref 6

CH3CH2OH CH3CHOH CH3COH CH2CH2OH CH2CHOH CH2COH CHCH2OH CHCHOH CHCOH CCH2OH CCHOH CCOH CH3CH2O CH3CHO CH3CO CH2CH2O CH2CHO CH2CO CHCH2O CHCHO CHCO CCH2O CCHO CCO H

−0.40 −0.61 −1.12 −0.29 −1.04 −1.33 −0.36 −1.05 −1.21 −1.22 −1.28 −0.25 0.12 −0.55 −1.39 −0.26 −0.61 −1.45 −0.24 −0.88 −1.57 −0.15 −1.09 −1.29 −0.62

−0.28 −0.29 −0.57 −0.01 −0.46 −0.72 0.01 −0.50 −0.65 −0.69 −0.63 0.38 0.39 −0.04 −0.87 0.14 −0.35 −0.85 0.31 −0.35 −0.91 0.86 −0.43 −0.55 −0.89

a

Excess H atoms are adsorbed on separate slabs with a binding energy calculated with respect to gas-phase H2.

binding energies from the work by Alcala et al.6 In general, the binding energies show good qualitative agreement. However, there is an approximately constant 0.5 eV offset between the two sets of binding energies, which could be attributable to differences in the computational setups. The earlier work uses 2 × 2 unit cells with 2-layer thick slabs frozen at the bulk position. While state-of-the-art at the time, computational resources have advanced significantly. The only C2 species (other than ethanol itself, which has nearly the same binding energy in both cases) which shows disagreement of more than about ±0.25 eV is CCH2O*. This is likely due to the differences in the optimized geometric structures, with ours being more stable (see the Supporting Information for the structure). This species is not expected to be important in either the most energetically favorable pathway or the microkinetic model. Additionally, we find that the least stable species from our calculations is ethoxy with a binding energy of 0.12 eV. Similarly, previous workers6 report a number of species with positive binding energies (i.e., species which are 4694



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Table 2. Comparison of Binding Energies (eV) of Ethanol Decomposition Products on Pt(111) with Respect to Gas-Phase Ethanol and with Additional H Atoms Adsorbed on Separate Slabs this work

ref 6

parent species

α C−H

β C−H

O−H

C−C

C−O

α C−H

β C−H

O−H

C−C

C−O

CH3CH2OH CH3CHOH CH3COH CH2CH2OH CH2CHOH CH2COH CHCH2OH CHCHOH CHCOH CCH2OH CCHOH CCOH CH3CH2O CH3CHO CH3CO CH2CH2O CH2CHO CH2CO CHCH2O CHCHO CHCO CCH2O CCHO CCO

−0.61 −1.12

−0.29 −1.04 −1.33 −0.36 −1.05 −1.21 −1.22 −1.28 −0.25

0.12 −0.55 −1.39 −0.26 −0.61 −1.45 −0.24 −0.88 −1.57 −0.15 −1.09 −1.29

−0.49 −0.82 −1.39 −0.42 −0.74 −1.31 −1.06 −1.39 −1.96 −0.59 −0.92 −1.49 −0.19 −0.97 −1.56 −0.12 −0.89 −1.48 −0.76 −1.53 −2.13 −0.29 −1.07 −1.66

0.01 0.03 −0.82 −0.30 −0.28 −0.48 −0.28 −0.26 0.55 −0.48 0.55 2.10 −0.14 −0.12 −0.97 −0.45 −0.43 −0.63 −0.43 −0.41 0.40 −0.63 0.40 1.95

−0.29 −0.57

−0.01 −0.46 −0.73 0.01 −0.50 −0.65 −0.69 −0.63 0.38

0.39 −0.04 −0.87 0.15 −0.35 −0.85 0.31 −0.35 −0.91 0.86 −0.44 −0.55

0.18 −0.12 −0.81 0.38 0.08 −0.60 −0.24 −0.55 −1.23 0.49 0.19 −0.50 0.33 −0.32 −1.08 0.54 −0.11 −0.87 −0.09 −0.75 −1.49 0.64 −0.01 −0.77

0.36 0.52 −0.39 0.17 0.23 0.19 0.23 0.36 1.46 0.19 1.46 2.83 0.62 0.78 −0.15 0.42 0.49 0.44 0.49 0.62 1.72 0.44 1.72 3.09

−1.04 −1.33 −1.05 −1.21 −1.28 −0.25 −0.55 −1.39 −0.61 −1.45 −0.88 −1.57 −1.09 −1.29

−0.26 −0.61 −1.45 −0.24 −0.88 −1.57 −0.15 −1.09 −1.29

less stable than gas-phase ethanol); one of these is ethoxy, which has the second most positive binding energy (their least stable species is CCH2O*). We consider five possible types (α C−H, β C−H, O−H, C− C, and C−O) of elementary decomposition reactions. The binding energies of the decomposition products are given in Table 2. As with the stable adsorbates, we find that our results agree qualitatively very well with the previous results.6 The binding energies of the dehydrogenation products follow almost the same distribution of deviations as the stable intermediates. This is because each C2 oxygenate produced via dehydrogenation is also a stable intermediate, leading to considerable overlap in the two sets of binding energies. In contrast, the binding energies of the C−C and C−O cracking products show slightly larger average deviations of around 0.8 eV. It has been previously reported38 that activation energies in the decomposition of ethylene glycol derivatives on Pt depend on the degree of hydrogenation. In this vein, we specifically examine the trends in binding energies and activation energies. We find that there is a general trend in the binding energies of the stable species with respect to the degree of dehydrogenation (Figure 1). Overall, it appears that the more highly dehydrogenated species bind more strongly. This trend is also observed in the previously published ethanol data.6 In fact, when comparing the binding energies, we find the exact same ordering of thermodynamic stability, as reported before for a given level of hydrogenation in every case but one, and in that case, the two species are nearly isoenergetic (less than 0.1 eV difference) such that distinguishing their relative stability via DFT is likely impossible. Next we turn our attention to the heats of reaction and activation energies. As might be expected from the results from

−0.46 −0.73 −0.50 −0.65 −0.63 0.38 −0.04 −0.87 −0.35 −0.85 −0.35 −0.91 −0.44 −0.55

0.15 −0.35 −0.85 0.31 −0.35 −0.91 0.86 −0.44 −0.55

Figure 1. Analysis of binding energies of stable surface intermediates with respect to gas-phase ethanol. Excess H atoms are adsorbed on separate slabs. Aside from adsorbed ethoxy, all surface intermediates are more stable than gas-phase ethanol. The binding energy generally increases with an increasing level of dehydrogenation. The weakly bound adsorbates generally correspond to highly strained species that are not expected to be kinetically relevant.

the binding energies, our heats of reaction show excellent qualitative agreement with prior results6 (data not shown). Our calculated activation energies along with the earlier literature values6 are given in Table 3. We also find excellent qualitative and sometimes quantitative agreement between our activation energies and the few reported in the earlier work. The principal difference is the earlier work did not calculate any barriers for dehydrogenation reactions. This lack of dehydrogenation 4695



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Table 3. Activation Energies (eV) for Decomposition Reactions in Ethanol Derivatives on Pt(111)a this work parent species CH3CH2OH CH3CHOH CH3COH CH2CH2OH CH2CHOH CH2COH CHCH2OH CHCHOH CHCOH CCH2OH CCHOH CCOH CH3CH2O CH3CHO CH3CO CH2CH2O CH2CHO CH2CO CHCH2O CHCHO CHCO CCH2O CCHO CCO a

α C−H 1.06 0.54 0.36 0.82

O−H

C−C

C−O

1.16 1.08 0.78 1.16 1.34 1.13

0.81 0.67 0.18 0.76 0.80 0.52 0.57 0.87 0.61 1.56 0.44 1.13

2.81 1.42 1.54 1.42 1.89 2.04 1.24 1.54 1.20 1.83 1.79 1.03 1.64 1.45 1.40 1.97 1.00 1.37 0.96 0.82 0.94 0.83 1.04 0.96

1.98 1.90 1.15 1.93 1.82 1.46 1.65 1.07 2.56 2.02 2.92 3.38 1.54 2.34 2.16 1.92 2.07 2.36 2.23 1.88 3.26 1.17 2.66 4.95

0.54 0.85

0.14 0.14 −0.01 0.44

ref 6

β C−H

0.64 1.27 0.98 1.33 1.06 0.79

0.57 0.50

C−C

C−O

1.01 2.00 1.63 1.40

1.46 1.95 1.85

1.23 1.34 1.48 1.36 1.38

2.27

0.95

1.42

No dehydrogenation barriers were reported in ref 6.

Figure 2. (a) Summary of dehydrogenation barriers vs the degree of dehydrogenation. There are no clear trends in the height of the barrier as a function of the level of dehydrogenation of the molecule. α C−H reactions have the lowest barriers, β C−H are the highest, and O−H are intermediate. (b) Summary of cracking barriers vs the degree of dehydrogenation. There is a clear trend in both the C−C and C−O(H) barriers. The C−C barriers decrease with increasing dehydrogenation, whereas the C−O(H) barriers generally increase. From the figures we can also draw some conclusions about the important activity and selectivity-controlling species. The lowest important C−C barrier (in CHCO) is around 1 eV, which is comparable to the initial α C−H scission in adsorbed ethanol. The lowest C−OH barrier in an important species is for CH3COH (around 1 eV), which is much higher than the lowest barrier in that species (O−H scission at 0.2 eV). These specific points suggest (1) that deoxygenation reactions are probably unimportant and (2) that C−C scission likely occurs late in the dehydrogenation sequence and may not be rate limiting.

C−H reactions tend to have the lowest barriers, β C−H reactions have the highest barriers, and O−H reactions are intermediate. C−C and C−O cracking reactions, however, do show trends. As was the case with ethylene glycol,38 we find that C−C barriers tend to decrease with the level of hydrogenation. C−O barriers, on the other hand, tend to increase with increasing levels of dehydrogenation.

barriers also has important implications for determining the kinetically important steps in a mechanism as discussed below. We also examine the activation energies for trends (Figure 2). We find that there is no clear trend in the dehydrogenation reactions. The level of dehydrogenation does not seem to have a significant effect on the barriers. Instead, the type of dehydrogenation reaction seems to play the biggest role; α 4696



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We attribute these trends to a combination of geometric and electronic effects. In the case of dehydrogenation reactions, with the possible exception of aromatic bonding, the level of hydrogenation does not strongly affect the geometric or electronic structure of the transition state. In contrast, the level of hydrogenation does affect the C−C and C−O bond structures. For C−C bonds, the more highly dehydrogenated species possess C−C bonds which are nearly parallel to the surface, reducing the barrier needed to break the bond. The opposite occurs with C−O bonds. As more H atoms are removed from the C−C−O backbone, the O tilts away from the surface. When this happens, in order to break the C−O bond, the CO(H) group must tilt back toward the surface before the C−O bond can break. This rotation is energetically disfavored and contributes to the increased barrier. In addition to investigating general trends in the activation energies, we can also use the data in Figure 2 to identify which species will be important in controlling the activity and selectivity of the thermal decomposition. We find that there are three species that have the potential to control the activity and selectivity. These are ethanol, 1-hydroxyethylidene (CH3COH*), and ketenyl (CHCO*). The dehydrogenation barriers in ethanol are similar to the C−C scission barrier in ketenyl, suggesting both that C−C scission probably occurs late in the process and that it may not be rate limiting. The relative difference and absolute magnitudes of the barriers in 1hydroxyethylidene suggest that O−H scission is preferred to dehydroxylation and that any deoxygenation will occur only in species with intermediate degrees of dehydrogenation (most probably in 1-hydroxyethylidene). We revisit this point in greater detail when we construct our most energetically favorable pathway. Oxidative Dehydrogenation Reactions. As with the thermal decomposition reactions, we can also calculate the heats of reaction and activation energies of the oxidative dehydrogenation reactions. Specifically, we consider six types of reactions: α C−H, β C−H, and O−H abstractions by both O* and OH* to form OH* and H2O*, respectively, together with the corresponding dehydrogenation product. We have calculated the aforementioned quantities both when reactants and products are coadsorbed on a single slab and when they are adsorbed on separate slabs. The binding energies for reactants and products in both reference states are given in Tables 4 and 5. The activation energies in both reference states are given in Table 6. Similar to the stable intermediates for the thermal decomposition reactions, there is a correlation of the binding energies with the degree of dehydrogenation. Specifically, the binding energy increases with the level of dehydrogenation (data not shown). More interesting, however, is the distribution of interaction energies (where the interaction energy is the difference between the binding energy in the coadsorbed state and the binding energy of the separate fragments). The distributions in the binding energies of both reactants and products are shown in Figure 3. In general, we find that species coadsorbed with atomic O tend to show a net repulsion, while species coadsorbed with OH* or H2O* (the latter from the product of OH*-assisted dehydrogenation) tend to either be slightly stabilized (OH*) or to show a wide variation (H2O*). One important exception to this is the OH* assisted dehydrogenation of terminal OH groups. In this case, there is a net stabilization of both the reactants and the products due to the ease of forming hydrogen bonds between the reactant OH*

Table 4. Binding Energies (eV) of Ethanol Derivatives on Pt(111) with Respect to Gas-Phase Ethanola α C−H

β C−H

O−H

species

separate

O

OH

O

OH

O

OH

CH3CH2OH CH3CHOH CH3COH CH2CH2OH CH2CHOH CH2COH CHCH2OH CHCHOH CH3CH2O CH3CHO CH2CH2O CH2CHO CH2CO CHCH2O CHCHO

−0.45 −0.67 −1.22 −0.47 −1.14 −1.50 −0.61 −1.22 0.17 −0.57 −0.29 −0.63 −1.52 −0.29 −0.97

0.13 −0.15

−0.44 −0.76

−0.29 −0.60

−0.56 −0.77

−0.39 −0.59

−0.52 −1.22

−0.03 −0.53 0.15 −0.53 −0.11 −0.54

−0.64 −1.19 0.28 −0.61 −0.29 −0.71

0.09 −0.56 −1.39 −0.52

−0.66 −1.22 −1.56 −0.61

−0.46 −0.68 −1.27 −0.01 −1.20 −1.38 −0.48 −1.15

−0.98 −1.27 −1.20 −1.10 −1.61 −2.20 −1.27 −1.62

0.28 0.06

−0.26 −0.53

0.28 −0.53 0.29 0.05 −1.53 −0.25

0.27 −1.57 −0.33 −0.67 −1.61 −0.54

a

Reactants are on separate slabs in the column indicated; otherwise they are coadsorbed on the same slab. Additional H atoms are adsorbed on separate slabs. All calculations are on 3 × 3 cells with a single additional O or OH (as indicated). The binding energies of different reaction types differ due to the placement of the coadsorbed O(H) species on the slab.

(product H2O*) with the terminal OH* (O*) of the ethanol derivative. Such stabilizing hydrogen bonds have been reported for ethanol on Rh.50,51 Additionally, we can compare the activation energies of the oxidative dehydrogenation reactions with the corresponding thermal dehydrogenation reactions. In order to easily compare trends in the activation energies, we examine the resulting distributions rather than individual reactions. These distributions are plotted in Figure 4. Starting with the α C−H reactions (Figure 4a), we find that there is little or no reduction in the dehydrogenation barriers when in the coadsorbed reference state. If we account for the reference state, the O* assisted reactions have actually increased barriers owing to the repulsive interaction between the coadsorbed reactants (i.e., the repulsive interaction is added to the barrier when in the separate slab reference). Similarly, there is a small reduction in the barrier for the OH assisted reaction due to a slight attractive interaction between the reactants. Similar results are also obtained for the β C−H reactions. Compared to thermal reactions, the O* (OH*)-assisted reactions have higher (lower) barriers (Figure 4b). After accounting for the slab reference state, the apparent barrier for the O* (OH*)-assisted reaction is further increased (decreased). Finally, with the O−H reactions (Figure 4c), we find that the abstraction by O* and OH* is favored in both cases over the thermal reactions due to the stabilization of the transition state by hydrogen bonding between the reactants. Interestingly, the O−H reactions with coadsorbed OH* are almost all nearly barrierless. This is supported by the structures of the stable reactants and products; in both states, the hydrogen atom being transferred is already very close to its transition state geometry (see the structures in the Supporting Information). Together with the large stabilization of the products and reactants via hydrogen bonding, this leads to apparently negative reaction barriers in both directions when the barrier is calculated with respect to the separate slabs. 4697



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Table 5. Binding Energies (eV) of Ethanol Dehydrogenation Products on Pt(111) with Respect to Gas-Phase Ethanol, with Products Adsorbed on the Same (Separate) Slab(s)a α C−H

β C−H

O−H

parent species

O

OH

O

OH

O


−0.82 (−0.57) −0.71 (−1.12)

−1.41 (−1.23) −1.14 (−1.78)

−0.46 (−0.37) −1.06 (−1.04)

−1.14 (−1.02) −1.55 (−1.70)

−1.11 (−1.04) −1.08 (−1.40)

−1.69 (−1.70) −1.99 (−2.05) −1.64 −1.89 −1.07 −1.82 −1.32 −2.38

(−0.53) (−1.12) (−1.25) (−1.18)

−1.29 (−1.16) −1.71 (−1.78) −1.80 (−1.91)

−1.02 −1.22 −0.44 −1.16 −0.75 −1.42

−0.53 −1.16 −1.24 −1.00

0.14 (0.27) −0.51 (−0.47) −1.18 (−1.34) −0.35 (−0.19) −0.59 (−0.53) −1.72 (−1.42) −0.41 (−0.20) −0.59 (−0.87)

0.06 (−0.19) −0.61 (−0.53) −0.15 (−0.20) −0.34 (−0.87) −1.50 (−1.56) 0.27 (−0.19)

−0.71 (−0.84) −1.45 (−1.19) −0.70 (−0.85)

(−1.12) (−1.25) (−0.47) (−1.34) (−0.53) (−1.42)

−0.35 (−0.87) −1.53 (−1.56)

(−1.78) (−1.91) (−1.12) (−1.99) (−1.19) (−2.08)

−1.12 (−1.53)

OH −0.87 −1.25 −2.03 −1.25 −1.57 −2.42 −1.28 −1.52

(−0.39) (−1.12) (−1.99) (−0.84) (−1.19) (−2.08) (−0.85) (−1.53)

−2.17 (−2.21) −0.61 (−0.85)

Additional H atoms are adsorbed on separate slabs. All calculations are on 3 × 3 cells with a single additional reactant O or OH (as indicated). The binding energies of different reaction types differ due to the placement of the coadsorbed reactant and product species on the slab. The heat of a particular dehydrogenation reaction may be calculated by subtracting the reactant binding energy (Table 4) from the corresponding product binding energy (this table). a

reaction and activation energies are calculated with respect to coadsorbed reactants and products and when they are calculated with respect to reactants and products adsorbed on separate slabs. This suggests that the use of coadsorbed reactants and products may not always be necessary; only in the cases where the interaction energy is significant (as it is for the O−H + OH reactions) this is important. Importantly, the amount of error due to this approximation can be quantified via recently developed methods.53 Such errors are likely to be small, with the intrinsic error of the BEP correlation dominating the effect of coadsorption in all but the most extreme cases. We can also compare the oxidative dehydrogenation parameters to previously published BEP parameters for thermal decomposition reactions of ethanol on Pt.52 Overall, we find that the parameters are very similar. This suggests that, in the absence of high concentrations of adsorbed O* or OH*, thermal decomposition pathways likely dominate. This was indeed found to be the case for ethylene glycol steam reforming on Pt.30 Most Energetically Favorable Pathway. With the binding energies, heats of reaction, and activation energies, we can assemble a most energetically favorable pathway. Given the expected unimportance of the oxidative dehydrogenation reactions, these are excluded from the most energetically favorable pathway. We present, then, a simplified scheme with the principal thermal decomposition pathways in Figure 6 for the reader to follow. Ethanol starts by adsorbing onto the surface from the gasphase with a binding energy of −0.40 eV. It can then undergo one of five possible reactions. Only three of these have low enough barriers to be kinetically relevant, α C−H scission, β C−H scission, and O−H scission, with barriers (reaction energy) of 1.06 (−0.21), 1.16 (0.12), and 0.81 (0.53) eV, respectively. Of these three, the first and the last are the most relevant. The α C−H scission is thermodynamically favored, while the O−H scission is kinetically favored. Thus, we follow both of these paths.

Table 6. Activation Energies (eV) of Oxidative Dehydrogenation Reactions on Pt(111)a α C−H parent species

O

OH


0.79 0.24

β C−H

O−H

O

OH

O

OH

0.89 0.43

1.35 1.06

0.80 0.70

0.61 0.64

0.40 0.62

0.88 0.09 0.43 0.58 0.27 1.11

0.72 0.27 −0.18 0.11 0.04 0.26

0.78 0.88 1.28 1.19

0.74 0.48 0.44 0.82

0.38 0.26 0.79 −0.15 0.56 −0.13 −0.03 0.55

0.06 −0.13 1.17 0.05 0.05 0.01 0.05 0.04

0.31 1.57 0.85

0.56 −0.06

0.55

2.06 1.34 0.43 1.66 1.02 2.12

0.52 1.62

a

All reactions are written in the dehydrogenation direction with reactants coadsorbed on the same slab. Activation energies with respect to separate slabs may be calculated from the difference between coadsorbed and separate slab reactant binding energies (both types of binding energies are found in Table 4).

Finally, we can use the reaction energy and activation energy data to develop Brønsted−Evans−Polanyi (BEP) correlations for the oxidative dehydrogenation reactions. Following the protocol outlined in our previous work,52 we can identify four statistically distinct homologous series: C−H abstraction by O* and OH* (denoted C−H + O and C−H + OH, respectively) and O−H abstraction by O* and OH* (denoted O−H + O and O−H + OH, respectively). The BEP correlations are shown in Figure 5, while the coefficients are given in Table 7. From Table 7, we can see that, with the sole exception of the O−H + OH reactions, at 95% confidence level there is no statistical difference in the coefficients when the heats of 4698



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Figure 3. Interaction energy distributions for stable reactants (a) and products (b) of oxidative dehydrogenation reactions of ethanol derivatives on Pt(111). Each box contains the middle 50% of the data, the intervening line represents the median, the whiskers represent either the actual minimum (maximum) or 1.5 times the interquartile range below (above) the first (third) quartile, and the circles represent points outside the whiskers that are deemed to be outliers. With the exception of species associated with the hydroxyl-assisted dehydrogenation of terminal hydroxyl groups, there is a repulsive or neutral interaction between coadsorbed species.

Figure 4. Comparison of activation energies for oxidative dehydrogenation of ethanol derivatives. Reactive centers are (a) α C, (b) β C, and (c) O. Reaction types are thermal decomposition (TD) and with O or OH on a coadsorbed slab basis. (Separate slab barriers are omitted for clarity; these may be trivially obtained by accounting for the interaction energy between the coadsorbed reactants.). Overall, the α C−H oxidative dehydrogenation barriers are all very similar to the thermal decomposition barriers. The β C−H oxidative dehydrogenation barriers show more variability with reactions with O (OH) having higher (lower) average barriers than thermal decomposition. In the O−H reactions, both reactions with O and OH are more favorable than simple thermal decomposition owing to the stabilization of the transition state by hydrogen bonding between the terminal OH and the reaction O(H) coadsorbate. The large outlier in panel (c), O−H + OH data set is the dehydrogenation of CH3COH*.

Table 7. Brønsted−Evans−Polanyi Correlation Parameters with 95% Confidence Intervals for Oxidative Dehydrogenation Reactionsa coadsorbed correlation

slope

intercept (eV)

C−H + O C−H + OH O−H + Ob

0.94 ± 0.40 0.68 ± 0.34 0.65 ± 0.63

1.34 ± 0.26 1.10 ± 0.30 0.16 ± 0.26

O−H + OHb

0.02 ± 0.59

0.02 ± 0.07

separate slope 0.61 ± 0.60 0.44 ± 0.43 −0.07 ± 0.94 0.41 ± 0.52

intercept (eV) 1.37 ± 0.24 0.90 ± 0.42 0.38 ± 0.38 −0.44 ± 0.18

a

There is no statistical difference between the coadsorbed and separate slab reference states for any of the correlations except for the O−H + OH reactions. bDehydrogenation of CH3COH excluded from correlation due to anomalously high activation energy. Figure 5. Brønsted−Evans−Polanyi correlations for distinct homologous series of oxidative dehydrogenation reactions. The heats of reaction are calculated using products and reactants coadsorbed on the same slab. Aside from the O−H abstraction by OH, at 95% confidence, there is no statistical difference in BEP parameters for the coadsorbed and separate slab cases.

Following the initial H abstraction, there are two possible parallel paths which ultimately reconverge further in the sequence. We start with the O−H bond scission to form ethoxy (CH3CH2O*) as the kinetically favored pathway. Out of the possible reactions, the α C−H scission to form acetaldehyde (CH3CHO*) is by far the most favorable, with a barrier 4699



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Figure 6. Scheme showing the most energetically favorable pathway in the decomposition of ethanol on Pt(111). Activation barriers are given first with the reaction energy following in parentheses. All energies are in eV. Ethanol, acetaldehyde, and CO undergo nonactivated adsorption, while methane adsorption is activated. The two key intermediates are adsorbed ethanol and acetyl (CH3CO*). The initial α C−H scission in adsorbed ethanol is the kinetically controlling step for total conversion, while the C−C scission in acetyl controls the selectivity to acetaldehyde vs methane and CO. Excess H* atoms are omitted for clarity.

Figure 7. Comparison of model (lines) and experimental (symbols) conversion (a) and selectivity (b). Experimental conditions are 12.5% ethanol, 37.5% H2O, bal. He at 1 atm total pressure. Experimental data were previously reported by Basagiannis et al.40 Overall, there is good agreement between the model and the experiments. The discrepancy in the COx and acetaldehyde selectivities at lower temperatures likely arises from neglecting the support. Ethane is also a minor product (maximum selectivity of around 2%) in both the model and in the experiments.

(reaction energy) of 0.14 (−0.68) eV. Acetaldehyde then can either desorb with a binding energy of −0.54 eV with respect to gas phase acetaldehyde or further decompose to acetyl (CH3CO*) via α C−H scission. This latter process has a barrier (reaction energy) of 0.14 (−0.84) eV. Thus, decomposition to acetyl is kinetically and thermodynamically favored. It is at acetyl that the parallel decomposition processes reconverge. Thus, we now turn our attention back to the initial α C−H scission in ethanol to form 1-hydroxyethyl (CH3CHOH*). This reaction is thermodynamically much more favorable (by 0.74 eV) than the competing process to form ethoxy. Coupled with the lower stability of ethoxy compared to gas-phase ethanol and the marginally higher barrier (only about 0.2 eV difference), it is anticipated that the reaction to form 1-hydroxyethyl will be the dominant pathway. Following its formation, 1-hydroxyethyl will decompose via two dominant pathways, to acetaldehyde with a barrier (reaction energy) of 0.67 (0.06) eV and to 1-hydroxyethylidene (CH3COH*) with a barrier (reaction energy) of 0.54 (−0.51) eV. The latter decomposition is kinetically and thermodynamically favored. From there, 1-hydroxyethylidene can either decompose into acetyl with a barrier (reaction energy) of 0.18 (−0.27) eV, or alternatively, it can undergo C−

OH scission with a barrier (reaction energy) of 1.15 (0.30) eV to form ethylidyne (CH3C*). The dehydrogenation to acetyl is strongly preferred kinetically and thermodynamically. Although dehydrogenation is strongly favored, the dehydroxylation barrier is low enough that minor quantities of C2 hydrocarbons could be formed by hydrogenating the resulting ethylidyne (see ref 35 for the C2 hydrocarbon pathway). From acetyl, the decomposition proceeds via one of two reactions. It can undergo β C−H scission with a barrier (reaction energy) of 0.98 eV (−0.05) to form ketene (CH2CO*) or undergo C−C cleavage with a barrier (reaction energy) of 1.40 (−0.17) eV to form methyl and CO*. The former process is strongly kinetically favored, while the latter process is slightly thermodynamically favored. The hydrogenation barrier to form acetaldehyde (0.98 eV) is identical to the dehydrogenation barrier to form ketene (CH2CO*), suggesting that the thermodynamically more stable ketene will be the preferred path. Ketene then preferentially dehydrogenates to ketenyl (CHCO*) with a barrier (reaction energy) of 0.79 (−0.12) eV. Ketenyl undergoes C−C cleavage with a barrier (reaction energy) of 0.94 (−0.56) eV to form methylidyne (CH*) and CO. All formed CHx* species undergo hydrogenation reactions to form methane (see ref 35 for the barriers and mechanism), while CO* desorbs unreacted. C−C 4700



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Figure 8. Apparent activation energies (a) and reaction orders at 300 °C (b) for the model and experiments for ethanol steam reforming. Experimental apparent activation energy data were previously reported by Basagiannis et al.40 Nominal experimental conditions are 12.5% (24%) ethanol, 37.5% (71%) water, bal. He at 1 atm total pressure for reaction orders (apparent activation energies). All model conversions are held constant at 5% by varying the model flow rate. Experimental conversions were held to less than 20% by varying the gas hourly space velocity. There is good agreement between the model and experimental apparent activation energies (a). Model and experimental reaction orders in ethanol (water) were obtained by varying the ethanol (water) partial pressure and holding the water (ethanol) partial pressure constant at its nominal value. Model (experimental) points are filled (open) symbols connected by solid (dashed) lines. Ethanol (water) points are represented by circles (squares). The rate is positive order with respect to ethanol and approximately zero order with respect to water, consistent with the overall paths being thermal decomposition followed by water−gas shift.

low amount of acetaldehyde indicates that this molecule is very likely to react rapidly on the catalyst surface producing H2, CO, and CH4 via ethanol dehydrogenation and acetaldehyde decomposition reactions. Our model is capable of quantitatively estimating the conversion under integral conditions (Figure 7a) with the tuning of a single extensive reactor parameter (the surface area to volume ratio, a measure of the amount of catalyst in the simulated reactor). The amount of tuning is within an order of magnitude of the estimated value for the experimental system, which is within the range of uncertainty typically associated with microkinetic models. Just as importantly, our model captures the correct qualitative trends for the selectivities (Figure 7b). There are quantitative deviations in the acetaldehyde, CO, and CO2 profiles, which are probably due to our modeling the metal chemistry only (i.e., neglecting the contribution of the support). Specifically, the support probably contributes to (1) the dehydrogenation to acetaldehyde, (2) the production of OH from water, and (3) small amounts of acid-catalyzed dehydration to ethylene (further hydrogenated to ethane by the Pt in our system) and condensation to diethyl ether (both species being observed in small amounts experimentally). Prior work by Basagiannis et al.40 showed that bare γ-Al2O3 was active for ethanol dehydrogenation and dehydration and proposed that it was also active for water activation. Nevertheless, experimentally we see very little dehydration or etherification. We believe that this is due to the preferential adsorption of water on the support at the expense of ethanol, which blocks acid sites and shifts the equilibrium toward ethanol. Such a hypothesis is supported by DFT calculations showing that water binds more strongly to the support than does ethanol.44 By analyzing the stoichiometric ratios of the products, it is possibly to identify which of the submechanisms, decomposition, water−gas shift, and methanation, are active. Some of the important limiting cases as well as the actual results from the microkinetic model are given in the Supporting Information. Of interest are the simple thermal decomposition

bond scission is preferred from highly dehydrogenated species and specifically from CHCO*. So it is possible the CH4 forms from scission of CHCO* followed by hydrogenation of CH*. Experimental and Microkinetic Modeling Results. Next we turn our attention to the microkinetic modeling of low-temperature ethanol steam reforming. We begin first by analyzing the conversion, selectivity, apparent activation energy, and reaction orders in both our model and our experiments to establish the validity of the model. We then interrogate our model results in order to gain mechanistic insights about the key reactions and intermediates and how they impact the ethanol steam reforming process. Our first test for the validity of our microkinetic model is simply comparing the model results for conversion and selectivity to the experimental results originally reported by Basagiannis et al.40 The model and experimental results are presented in Figure 7. Experimentally, ethanol conversion (Figure 7a) increases with increasing temperature and reaches 100% at around 360 °C. The main products detected are H2, CH4, CO, CO2, and to a smaller extent, at low temperatures, CH3CHO (Figure 7b). Selectivity toward hydrogen production increases from 38 to 47%, with increasing temperature from 300 to 400 °C, whereas selectivity toward methane is rather stable over the entire temperature range at the level of 44 to 47%. Carbon monoxide and carbon dioxide selectivities present a symmetric variation, implying production of CO2 at the expense of CO. In particular, the selectivity to CO decreases substantially above 320 °C from 38 to 7%, while the selectivity to CO2 varies inversely to that of CO, increasing from 9 to 44% in the temperature range of 300−400 °C. This behavior is due to the occurrence of the WGS reaction, in agreement with previous studies.40,41,54 The production of H2, CH4, and CO can be attributed to ethanol decomposition or dehydrogenation with concomitant adsorption or rapid decomposition of acetaldehyde. The latter pathway is supported by the low selectivity toward CH3CHO (less than 7%) observed in the whole temperature range examined (Figure 7b). However, the very 4701



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Figure 9. Reaction path analysis at the reactor exit for ethanol steam reforming at 300 °C and for feed conditions from Figure 7. Solid (dotted) lines represent primary (secondary) pathways. Excess H* atoms being added or removed are omitted for clarity. Reversible (irreversible) reactions are represented by double (single) headed arrows. Percentages are for the total amount of C in the ethanol which proceeds through a given irreversible reaction and may not add to exactly 100% due to rounding. Fractions denote the partial equilibrium index (PEI). The initial α dehydrogenation of ethanol is irreversible and occurs as fast as the net rate of adsorption. Subsequent dehydrogenations are partially equilibrated until the formation of ketenyl (CHCO*), which undergoes irreversible C−C bond scission as quickly as it is formed. This suggests that the initial dehydrogenation is the rate determining step for total conversion, while the C−C scission controls the selectivity to acetaldehyde. The production of CO2 is limited by the slow formation and decomposition of carboxyl (COOH*).

negative with the rate decreasing by a factor of 1.4. The slight negative order in water is consistent with water blocking sites on the support that are active for ethanol dehydrogenation and dehydration. Our reaction orders are consistent with the reaction orders for ethylene glycol steam reforming on Pt30 and with the other ethanol steam reforming studies cited earlier. As with our selectivity data, our reaction orders are consistent with a mechanism dominated by thermal decomposition, with water−gas shift playing a minor role in the overall rate. We begin our analysis of the mechanism by performing a reaction path analysis (RPA). With this analysis we identify the principal pathways in the mechanism. We do this by following the evolution of individual species through the simulated reactor as they participate in their associated elementary reactions. Three quantities are important for identifying the important reactions and species and diagramming the most important pathway: (1) the net rate of reaction, (2) the partial equilibrium index (PEI), and (3) the relative importance of each reaction in the formation or consumption of each species. The PEI is defined as

and thermal decomposition followed by water−gas shift mechanisms. The simple thermal decomposition mechanism has all ratios equal to unity. Likewise, the thermal decomposition followed by complete water−gas shift pathway still has a CH4/COx ratio of unity, but a H2/COx ratio of two (a CO/COx ratio of zero). Our model product ratios are intermediate between these two cases at all temperatures, with water−gas shift becoming more important at higher temperatures. At higher temperatures, the principal reactions in the water−gas shift mechanism approach equilibrium, suggesting that the CO/COx ratio is thermodynamically rather than kinetically controlled (see Supporting Information for a comparison of our compositions to the equilibrium compositions). This general trend is also seen experimentally. Overall, both our model and experimental selectivity results are consistent with an overall mechanism of simple thermal decomposition followed by varying (dependent on support) amounts of water−gas shift. Our second test is of the apparent activation energy (experimental data previously reported by Basagiannis et al.40) and reaction orders in ethanol and water at 300 °C. Comparisons between the experimental and model results are given in Figure 8. We find good qualitative agreement between our model and experimental apparent activation energies (Figure 8a). Our model of apparent activation energy is somewhat higher than that of the experimental data. Possible sources of this discrepancy include (1) a difference in the conversions at corresponding model and experimental data points (the model was held to 5% conversion everywhere, while the experimental conversions were somewhat higher), (2) the participation of the γ-Al2O3 support in the mechanism which would increase conversion of ethanol via dehydration to ethylene and possibly dehydrogenation to acetaldehyde, and (3) uncertainty in the model estimate and/or from experimental data. The model also indicates that the ethanol steam reforming reaction is nearly first-order with respect to ethanol and zero-order with respect to water (Figure 8b). Our experimental reaction orders are consistent with our model reaction orders, with ethanol being nearly first order, with the rate increasing by about a factor of 6 and water being slightly

PEI =

rfwd rfwd + rrev

(27)

where rfwd and rrev are the absolute values of the forward and reverse rates of reaction, respectively. It is a quantitative measure of how fast the forward and reverse reactions are with respect to each other. Any reaction with a PEI between 0.45 and 0.55 is deemed to be in partial equilibrium (i.e., fully reversible), while a larger (smaller) ratio indicates the reaction is irreversible in the forward (reverse) direction. The important reactions in the mechanism typically have relatively fast rates and are irreversible. Other reactions will either not occur (if the net rate is too small) or will have no effect on the conversion (if they are in partial equilibrium). By chaining together each of the important reactions we arrive at a scheme showing the important path(s) from reactants to products. Although the pathways may change as a function of conversion (i.e., as a function of location in the reactor), the changes for our model are relatively minor and are primarily in the secondary pathways. The resulting scheme for our model at 300 °C and 4702



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Figure 10. Coverages of important surface intermediates at the reactor exit and for feed conditions from Figure 7. The principal decomposition products are present in relatively high coverages at all temperatures (panel a), while other key intermediates in the principal pathway are present in lesser amounts (panel b). The coverage profiles for 1-hydroxyethyl (CH3CHOH*), 1-hydroxyethylidene (CH3COH*), and ketene (CH2CO*) are excluded for clarity; 1-hydroxyethyl roughly follows the profile of ethanol, while 1-hydroxyethylidene and ketene roughly follow the profile of ketenyl. The similarities of these profiles are consistent with most of the dehydrogenation reactions (except the early ones) being in partial equilibrium. The lower coverages of the key intermediates are responsible for the low reaction rates of key reactions, especially the initial dehydrogenation of ethanol, the C−C cracking of CHCO*, and the formation of COOH* (and ultimately CO2) from CO* and OH*.

0.69 for the dehydrogenation of CH3*), while CO2 formation is too slow to affect the rate (see the Supporting Information, Figure S39, for a plot of selected gas-phase mass fractions illustrating that the thermal decomposition is nearly complete before water−gas shift starts). Overall, this suggests that either the initial α C−H abstraction or the C−C cracking of ketenyl will be the rate controlling step. The RPA also suggests that, as the principal C−C cracking step, the decomposition of ketenyl controls the selectivity to acetaldehyde versus the formation of C1 species. We also find that other overall reactions included in the model, for example, the hydrogenation of CO/CO2, the steam reforming of CH4, and the formation of C2 hydrocarbons, are of negligible importance at these conditions. For the next part of our analysis, we examine the coverages of species on the surface. These are presented in Figure 10. The surface coverages can give us further insight into which species and reactions are important. For convenience, we have divided the species into two categories: principal decomposition products (panel a) and reactive intermediates (panel b). We begin with the principal decomposition products. By far the most abundant “species” is surface vacancies, which indicates that blocking of sites will not significantly inhibit the rate, a postulate which is supported by the positive reaction order in ethanol (0.82 from the experiments and 0.68 from the model) from Figure 8. It is also readily apparent that the principal decomposition products CO* and H* are relatively abundant on the surface. The final principal decomposition product methylidyne is less abundant, primarily because it is also an important reactive intermediate in the pathway to methane formation. The relative abundance of these species further suggests that their presence does not control the rates of bimolecular reactions in which they participate as reactants (e.g., carboxyl formation and CHx* hydrogenation). Turning our attention to the reactive intermediates, we immediately note that the coverages of many of these species are much lower than the coverages of the principal decomposition products. It is clear that the most abundant species are not, in fact, the important reactive intermediates. Adsorbed OH is particularly scarce on the surface, and this

for the reactor outlet is given in Figure 9. Overall, this scheme corresponds very well with the DFT results in Figure 6, with the principal difference being the elimination of the formation of ethoxy. First, ethanol adsorbs reversibly on the surface. Then nearly 100% of the adsorbed ethanol (on a C basis) irreversibly decomposes via α C−H abstraction to form 1-hydroxyethyl (CH3CHOH*) with only minor amounts (less than 1%) reversibly forming ethoxy (CH3CH2O*) via O−H scission. The large disparity in selectivity between these two reactions is due to the thermodynamic stability of the reactants and products. The model shows that the O−H reaction to form ethoxy has an equilibrium constant much smaller than unity (order 10−6), while the α C−H reaction to form 1-hydroxyethyl is much larger than unity (order 100) resulting in the latter pathway being the dominant one. Subsequently, acetyl (CH3CO*) is formed primarily via the 1-hydroxyethylidene (CH3COH*) intermediate. Acetyl is also in partial equilibrium with adsorbed and gaseous acetaldehyde and other species. Very small amounts of acetyl undergo C−C scission to form CO* and CH3* (less than 1% of the original C each), but acetyl principally undergoes successive partially equilibrated dehydrogenations to form ketenyl (CHCO*). Ketenyl then cracks at the C−C bond to form CO* and methylidyne (CH*) in equal quantities (greater than 49% of the original C each). Minor amounts of CO* (containing less than 1% of the original C) react with OH* (formed principally from water, although some is also generated from minor amounts of C−O scission in CH3COH*) to ultimately form CO2 (the rest simply reversibly desorbs). Based on the reaction path analysis, we can draw some preliminary conclusions regarding the rate determining steps for conversion and selectivity to different products. Ethanol adsorption is partially equilibrated, but the subsequent dehydrogenation to 1-hydroxyethyl is irreversible. All subsequent dehydrogenation reactions to ketenyl are in partial equilibrium. Following the irreversible C−C cracking of ketenyl, subsequent steps for the formation of methane are either partially equilibrated or very nearly so (the highest PEI is 4703



Article

its NSC is less than the NSC of the initial dehydrogenation step by an order of magnitude. As a result, its most important role in the mechanism is in controlling the selectivity. Indeed, in sensitivity analysis of selectivity, we have confirmed that this reaction is the one with the highest impact on the production of acetaldehyde (data not shown). Other reactions are less sensitive than these two and have little effect on the conversion or selectivity. Had we assumed the initial dehydrogenation was equilibrated, we would have misidentified the rate determining step, and the model rates would have been orders of magnitude higher than our experimental rates.

dearth controls the rate of carboxyl (and ultimately CO2) formation. The very low coverage of OH* justifies our neglect of oxidative dehydrogenation reactions in assembling our model. The relatively high coverage of acetyl is due to its thermodynamic stability relative to acetaldehyde and ketene, but as all previous and subsequent hydrogenation/dehydrogenation reactions are in partial equilibrium, acetyl decomposition should not affect the rate. Methylene and methyl are also less abundant than methylidyne, implying that their hydrogenation to methyl is faster than the production of methylidyne from ketenyl cracking. However, of more importance for identifying the rate determining step are the coverages of ethanol and ketenyl. If we consider only the activation energies at 300 °C, ketenyl decomposition (21 kcal/mol barrier) should be rate controlling with a rate 2 orders of magnitude lower than ethanol dehydrogenation (15 kcal/mol barrier). However, at all temperatures adsorbed ethanol is about 3 orders of magnitude less abundant than adsorbed ketenyl. The net result is that the dehydrogenation of ethanol to 1-hydroxyethyl will be about an order of magnitude slower than the C−C cracking of ketenyl. Thus, the rate determining step for conversion should be the initial dehydrogenation of ethanol while the selectivity to acetaldehyde should be the C−C cracking reaction, contrary to the kinetic barrier prediction. We can confirm our identification of the rate determining step more rigorously via a sensitivity analysis. In a sensitivity analysis, we perturb a kinetic parameter (in our case the preexponential) in each reaction in the mechanism and calculate its effect on an output quantity (e.g., conversion or composition). With this information, we can calculate the normalized sensitivity coefficient (NSC) for the i-th reaction by NSCi =

d ln X d ln ki

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CONCLUSIONS In this work we performed a combined DFT, experimental, and microkinetic modeling study of low temperature ethanol steam reforming on Pt. Using DFT we mapped out the complete ethanol thermal decomposition pathway and explored oxidative dehydrogenation reactions as possible alternatives to pure thermal decomposition under steam reforming conditions. We examined trends in activation energies for both types of reactions. We used our DFT results to construct a most energetically favorable pathway for ethanol thermal decomposition. We also constructed the first comprehensive microkinetic model for low temperature ethanol steam reforming on Pt and compared it to suitable kinetic data. Finally, we used reaction path and sensitivity analyses along with the coverages of reactive intermediates to identify the rate determining step and important chemistry. Our DFT calculations show that there are discernible trends in the binding energies and some activation energies. Specifically, we find that with increasing levels of dehydrogenation (1) the binding energies increase, (2) the activation energies for C−C cracking reactions decrease, and (3) the activation energies for C−O and C−OH reactions increase. The dehydrogenation barriers are not strongly affected by the level of hydrogenation. While oxidative dehydrogenation can be energetically favorable compared to pure thermal decomposition reactions, their contribution is relatively small due to the low coverage of adsorbed OH. From our most energetically favorable pathway, we find that the initial α C−H abstraction is strongly thermodynamically favored over the O−H abstraction but that a comparison of activation energies is unable to rule out one pathway over the other. The reaction then proceeds via subsequent dehydrogenations to ketenyl, which undergoes C− C cracking to form CO* and methylidyne which eventually hydrogenates to form methane. The initial α C−H abstraction in ethanol and C−C cracking in ketenyl have similar barriers, making it unclear which reaction is the rate determining step. We then compared our microkinetic model results to kinetic data. We showed that our model is capable of (1) quantitatively predicting the experimental conversion (with the fitting of a single extensive reactor parameter, namely the surface area) and (2) qualitatively matching the important selectivity trends with quantitative deviations attributable primarily to the model’s neglect of the support. We found that the overall mechanism is satisfactorily described by simple thermal decomposition followed by water-gas shift. Finally, we used our microkinetic model to interrogate the mechanism of low temperature ethanol steam reforming. Our reaction path analysis showed that the mechanism largely follows the same sequence of reactions predicted by the DFT most energetically favorable pathway but excludes some energetically favorable DFT pathways. It also clarified that

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where X is the conversion and ki is the pre-exponential of the ith reaction. The NSC is a quantitative measure of the sensitivity of each reaction on the overall conversion with larger absolute values indicating higher sensitivity. The normalization ensures that all sensitivity coefficients are directly comparable. We present the results of our sensitivity analyses at 300 and 340 °C in Figure 11. We find that the initial α C−H abstraction is the most sensitive reaction at both temperatures. The C−C cracking of ketenyl is the second most sensitive reaction, but

Figure 11. Normalized sensitivity coefficients (see text for definition) at 300 and 340 °C for feed conditions from Figure 7. The most sensitive reaction is found to be the initial α C−H abstraction in adsorbed ethanol. Other reactions are much less sensitive and have little effect on the overall conversion. 4704



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ethoxy formation is unimportant on Pt due to the higher thermodynamic stability of the α C−H abstraction product 1hydroxyethyl and that C−C cracking occurs primarily in the ketenyl intermediate. Either the initial dehydrogenation or the C−C cracking of ketenyl could be the rate determining step with the C−C cracking controlling the selectivity to acetaldehyde. Analysis of the coverages of important surface intermediates confirmed the importance of these two particular reactions for controlling the rate and selectivity. The initial dehydrogenation was found to be the rate determining step because of the much lower concentration of adsorbed ethanol relative to ketenyl. The identification of the initial α C−H abstraction as the rate determining step, with the C−C cracking of ketenyl as the reaction controlling selectivity to acetaldehyde was confirmed via a sensitivity analysis. Our results clearly indicate that dehydrogenation/hydrogenation reactions may be more important than previously thought in the reforming of oxygenated molecules and that an assumption about equilibrated chemistry can lead to erroneous conclusions for the overall rate determining step in reforming chemistry.

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ASSOCIATED CONTENT

S Supporting Information *

Graphical representations of the optimized geometries of each DFT calculations, justification for our neglect of the partition functions in calculating the pre-exponentials, a full set of parameters for the microkinetic model, a discussion of the thermodynamic equilibrium of the water−gas shift reaction, tables of the product ratios showing that the product stoichiometry is consistent with thermal decomposition followed by incomplete water−gas shift, and a further discussion of the rate of thermal decomposition compared to the rate of water−gas shift. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel.: 302-831-2830. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors thank Dr. Ying Chen for assistance with the ethanol/Pt DFT calculations and Dr. Michael Salciccioli and Mr. Matthew Christiansen for helpful discussions. This work was partially supported by the NSF (Award No. CBET940768).

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REFERENCES

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Combined DFT, Microkinetic, and Experimental Study of Ethanol

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