Density Functional Theory and Microkinetic Studies of Bio-oil

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DFT and Microkinetic Studies of Bio-Oil Decomposition on a Cobalt Surface: Formic Acid as a Model Compound Xinbao Li, Shurong Wang, Yingying Zhu, Chen Lv, and Guohua Yang Energy Fuels, Just Accepted Manuscript • DOI: 10.1021/acs.energyfuels.6b03005 • Publication Date (Web): 07 Jan 2017 Downloaded from http://pubs.acs.org on January 15, 2017

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DFT and Microkinetic Studies of Bio-Oil Decomposition on a Cobalt Surface: Formic Acid as a Model Compound

Xinbao Li a, Shurong Wang b,*, Yingying Zhu a, Chen Lv a, Guohua Yang a * Corresponding author e-mail: [email protected] Tel.: +86 571 87952801; Fax: +86 571 87951616 a

Faculty of Maritime and Transportation, Ningbo University, Ningbo 315211, P R China

b

State Key Laboratory of Clean Energy Utilization, Zhejiang University, Hangzhou 310027, P R China

Abstract: Density functional theory calculations and microkinetic modeling were used to study the decomposition mechanisms of the bio-oil model compound formic acid over a cobalt-stepped surface. Zero-point energy-corrected activation barriers and reaction energies and the rate and equilibrium constants of various elementary reactions were obtained. Formic acid dissociation likely starts from dehydrogenation and dehydroxylation, with activation barriers of less than 0.5 eV. The generation of an HCOO intermediate is thermodynamically favored, but such a compound is energetically difficult to convert. COOH formation is fast and dominant at low temperatures, and it is converted rapidly after 450 K. The most favorable formic acid decomposition pathway is HCOOH → COOH → CO. Its rate-determining step is CO–OH scission, with an activation barrier of 0.66 eV and strong exothermicity of 1

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–1.19 eV. Keywords: DFT, microkinetic, bio-oil, formic acid, cobalt, hydrogen

1. Introduction Hydrogen is a clean energy source that can substitute for polluting fossil fuels. It is used extensively in fuel cells and combustion engines. Biomass-derived hydrogen production has gained increased attention globally because of its raw material renewability. Bio-oil catalytic reforming is a viable and competitive technology for biomass-derived hydrogen production.1,2 Liquid bio-oil can be produced via biomass fast pyrolysis with a high productivity of up to 75%.3 Its characteristic low cost, high productivity and convenient transportability make the bio-oil catalytic reforming technology more economical and easier to scale up. As a result, the technology has grown rapidly and has been studied extensively over the past two decades.4−8 However, limitations on catalyst activity and stability have become a core challenge for its commercial application.4,9 To improve the efficiency of the technology, a clear and fundamental understanding of the detailed reaction mechanism and accurate thermodynamic and kinetic parameters of the elementary reactions is required. Bio-oil is a complex oxygenated mixture of carboxylic acids, alcohols, ketones, aldehydes, phenolic derivatives, furans, sugars, and heavier oxygenates.10,11 Carboxylic acids are the main compounds of bio-oil, and these comprise mostly acetic and formic acids.12−14 Model compounds, especially acetic acid, have been used frequently in studies.6 We have carried out a series of mechanistic studies on bio-oil 2


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decomposition using acetic acid (CH3COOH) as a model compound over Pd(111), Ni(111), and Co(111) flat surfaces and a Co-stepped surface via density functional theory (DFT) calculations.7,12,15,16 To understand the reaction mechanism further, formic acid (HCOOH) was selected to perform DFT and microkinetic-combined calculations over a cobalt (Co) catalyst surface. Co-based catalysts have superior reactivity and stability in bio-oil catalytic reforming for hydrogen production.12,17 Xing et al.8 suggested that the Co catalyst may be a desirable economic alternative for the steam reforming of biomass-derived oxygenates compared with more conventional Ni and Rh-type steam reforming catalysts. HCOOH has also attracted attention as a promising hydrogen-carrier material. DFT calculations for its decomposition have been previously studied over various metal surfaces (Ag, Cu, Pd, Pt, Rh, Ni, and Au).18−21 Two parallel pathways involving dehydrogenation (HCOOH → H2 + CO2) or dehydration (HCOOH → H2O + CO) reactions have been identified for HCOOH decomposition. However, to the best of our knowledge, HCOOH decomposition on a Co-stepped surface has not been reported. It is well recognized that stepped surfaces are the most active sites for heterogeneous catalytic reactions.6,15,22 The purpose of this study was to gain insight into the energetics and kinetics of HCOOH decomposition. This work is organized as follows: In Section 2, the detailed computational methods for DFT calculations and microkinetic modeling are given. In Section 3, the zero-point energy (ZPE)-corrected activation barriers and reaction energies of each elementary reaction are presented. The forward rate constant and equilibrium constant 3


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are also presented. The most likely reaction pathway is discussed and identified using microkinetic modeling. Conclusions are given in Section 4. 2. Computational Methods 2.1. DFT calculation Spin-polarized periodic DFT calculations were performed using the Vienna Ab initio Simulation Package (VASP).23−25 Exchange and correlation effects were calculated using Perdew-Burke-Ernzerhof (PBE) generalized gradient correction (GGA). A projector-augmented wave pseudopotential was applied to describe the core electrons of all atoms. Monkhorst-Pack k-point meshes of 6×6×6 and 3×3×1 for bulk and surface calculations were used to sample the Brillouin zones. The kinetic energy cut-off was set to 400 eV, and the Fermi level was smeared using the Methfessel-Paxton approach with a width of 0.1 eV. The electronic self-consistent field was converged to 1×10−4 eV, and forces on all atoms were converged to less than 0.05 eV/Å. The optimized lattice constant for bulk fcc Co is 3.524 Å, which agrees well with theoretical (3.521 Å) and experimental values (3.545 Å).26,27 The construction of a Co-stepped surface was described in our previous study.15 Transition state (TS) searching and minimum energy path (MEP) finding were performed by the climbing nudged elastic-band method (CINEB).28,29 For the frequency calculations, all the metal atoms were fixed, and the adsorbates were free to relax in all directions. A ZPE correction was used for all adsorption energies, reaction energies, and activation barriers. 2.2. Microkinetic modeling 4


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We applied a harmonic transition-state theory (hTST) to calculate the rate constant of the elementary reactions. The forward rate constant (kf) of each reaction was calculated via the Eyring equation: 3N

kT kf = B h



 hviIS  

∏ 1 − exp  − k T 

Eac

 B   − k BT  e 3 N −1   hviTS   1 − exp  −  ∏ i =1   k BT   i =1

where kB is the Boltzmann constant, h is the Planck constant, T is the absolute temperature, vi is the vibrational frequency of each vibrational mode of the adsorbed intermediate, and Eac is the ZPE-corrected activation energy. The reverse rate constant (kr) is calculated similarly, and the thermodynamic equilibrium constant Keq is defined as:

Keq =

kf kr

According to the procedure of Lu et al.,30 we shifted all real frequencies below 200 cm–1 to 200 cm–1 to cancel out the error caused by the low frequency. It is thought that low-frequency modes have no effect on reaction energies and barriers for surface reactions. We simplified and assumed that (i) all intermediates were located at the surface, (ii) no adsorption/desorption reactions occurred, and (iii) all elementary reactions were reversible, except for H2 and H2O formations. Modeling was carried out in an isothermal steady-state plug-flow reactor. 3. Results and Discussion 3.1. Elementary Reactions We proposed four different pathways for HCOOH decomposition, namely, HCOO-, 5


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COOH-, HCO-, and HCOH-mediated pathways, as shown in Fig. 1. The corresponding ZPE-corrected activation energies, reaction energies, and rate parameters for all elementary reactions are listed in Table 1. 3.1.1. HCOOH initial dissociation HCOOH dissociation may start from R1) O–H bond cleavage to produce HCOO, R2) C–OH bond cleavage to produce HCO, R3) C–H bond cleavage to produce COOH, and R4) carbonyl C=O bond cleavage to produce HCOH. The corresponding reaction barriers and potential energy profiles are presented in Table 1 and Fig. 2. The activation barriers of these four reactions are lower than 1 eV, which implies the easy dissociation of HCOOH over a Co surface. The reaction with the highest activation barrier is R4 (HCOOH → HCOH + O, Eac = 0.79 eV), while that with the lowest barrier is R3 (HCOOH → COOH + H, Eac = 0.26 eV). Therefore, HCOOH dissociation is kinetically preferred, starting from C–H cleavage. This is consistent with our previous study for acetic acid decomposition over a Co-stepped surface, in which the α–carbon dehydrogenation of CH3COOH was also identified as the lowest-energetic starting step associated with a close activation barrier of 0.32 eV.15 The barriers for R1 and R2 are only 0.10 and 0.19 eV higher than that of R3, which indicates that energetically similar dissociation reactions exist where HCOOH undergoes dehydrogenation to formate (HCOO) and dehydroxylation to formyl (HCO). The obtained barriers are also close to our previous results of CH3COO–H scission and CH3CO–OH scission.15 All four reactions (R1–R4) are exothermic. R1 has a maximum exothermicity of –1.12 eV, whereas R3 is mildly exothermic (–0.42 6


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eV). R4 is slightly exothermic (–0.19 eV). Yoo et al.18 reported that the barriers for O–H and C–H scissions of HCOOH over Pt(111) were 0.30 and 0.22 eV, respectively, which are close to our results. The O–H scission over Ni(111) computed by Luo et al. also showed a similar barrier of 0.41 eV.19 While Singh et al.20 found that the C–H scission over Au(111), Au(100), and Au(211) have barriers of 1.22, 0.83, and 0.87 eV, respectively, which are much higher than our result. Co reveals a comparable reactivity with the superior Pt catalyst in HCOOH initial decomposition. The forward rate constants (kf) and equilibrium constants (Keq) of R1–R4 are depicted in Fig. 3. The values calculated at 773 K are listed in Table 1. Figure 3(A) shows that the reaction rate is in the following order: R3 > R1 > R2 >> R4. At 773 K, as an example, the corresponding rate constants are 1.78×1011, 1.03×1011, 1.46×1010, and 5.96×107 s–1, respectively. Therefore, R3 is confirmed to be the most kinetically preferred reaction for initial HCOOH dissociation. The kf of R1 approaches R3 with an increase in temperature, especially above 850 K. It can be speculated that at a high temperature, R1 is competitive with R3 because of their similar reaction rates. Figure 3(B) shows that Keq varies as R1 >> R2 > R3 > R4. At 773 K, their values are 1.27×107, 2.19×104, 6.33×102, and 2.06×101, respectively. R1 is almost 5 orders of magnitude higher than R3, which suggests that R1 is the most thermodynamically preferred step for initial HCOOH dissociation. The Keq of R1–R4 decrease with an increase in temperature, which implies that the forward reactions of the HCOOH first-step decomposition will benefit from a decrease in reaction temperature. This is 7


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consistent with the fact that exothermic reactions (R1–R4) are favored by a decrease in temperature. Snapshots of the initial state (IS), transition state (TS), and final state (FS) of R1–R4 are shown in Fig. 4. As a reactant, HCOOH is initially adsorbed on two stepped Co atoms (sCo) through the C and carbonyl O atoms to form a η2–µ2 (C,O) configuration, and it has a mild adsorption energy of –0.96 eV. The corresponding distance of d(C–sCo) is 2.01 Å, and d(O–sCo) are 1.98 and 1.96 Å. For R1, the O–H bond in TS1 is elongated from 0.99 Å to 1.34 Å to give a co-adsorbed HCOO and H (FS1). HCOO forms a bridge configuration with two O atoms on top of the sCo atoms. The associated d(O–sCo) are both 1.94 Å. H locates on a hollow site of the flat surface. The co-adsorbed HCO and OH in the FS2 is a product of HCOOH dehydroxylation (R2). The C–OH distance in TS2 is elongated from 1.36 to 2.02 Å. For dehydrogenation (R3), the produced H binds to two sCo atoms on the step, and the co-produced COOH adsorbs to the surface via its unsaturated C atom on a bridge site and its carbonyl O atom on the top site. The length of C–H in TS3 is nearly identical to that in IS, but the angle between C–H and the surface is tilted from 53° to 32°, which is associated with the O–H being rotated from 59° to 24°. The C–H and O–H point toward and away from the surface, respectively. For R4, the cleaving C=O bond in TS4 is 0.59 Å longer than that in IS. 3.1.2. HCOO and HCOH dissociations As the most thermodynamically preferred product of initial HCOOH dissociation, 8


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HCOO may undergo dehydrogenation to yield CO2 (R6), and/or deoxidation followed by dehydrogenation to form CO (R5, R8). The potential energy profiles are presented in Fig. 5. R6 (HCOO → CO2 + H) and R5 (HCOO → HCO + O) have similar activation barriers and reaction energies. The barrier of R6 is 1.33 eV, which is only 0.03 eV higher than that of R5. The dissociation of HCOO is more difficult than its formation from HCOOH because the activation barriers of R5 and R6 are ~1.0 eV higher than that of R1, which leads to a possible accumulation of HCOO on the surface. R5 and R6 are mildly endothermic (0.66 and 0.59 eV), respectively. Scaranto et al.31,32 reported that the barriers for the dehydrogenation of HCOO over Pt(111) and Pd(111) are approximately 1 eV, and those for the deoxidation over Pt(111) and Pd(111) are higher than 1.8 eV. Yoo et al.18 found that the barriers for the dehydrogenation of HCOO over Pt(211), Pd(211), Rh(211), Cu(211) and Ag(211) stepped surfaces ranged from 1.00 to 1.45 eV. These published results suggest that the dissociation of HCOO on metal surface is difficult, which is in good agreement with our calculation. The subsequent dehydrogenation of HCO to CO in R8 has an extremely low activation barrier of 0.03 eV, which is associated with a strong exothermicity of –0.97 eV. The ZPE correction was not applied for this activation barrier because its value is almost zero. This approximation will not influence the global reaction pathways and reaction rate significantly because it is a tiny barrier. As a result, the intermediate of HCO can be readily converted to CO over the Co surface. Another intermediate of HCOH, which is generated from R4, can dissociate to HCO via HCO–H scission with 9


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an activation barrier of 0.61 eV and a mild exothermicity of –0.42 eV (R7). R8 then follows to yield CO. The corresponding kf and Keq of R5–R8 are shown in Fig. 6. The kf increases as R8 >> R7 >> R6 ≈ R5. At 773 K, the values are 4.99×1012, 1.11×109, 3.97×104, and 3.45×104 s–1, respectively (Table 1). The kf of R8 remains nearly constant with changes in temperature because of its negligible activation barrier. In comparison with HCOO formation (R1 with kf1 = 1.03×1011 s–1), its decomposition rates (R6 with kf6 = 3.97×104 s–1 and R5 with kf5 = 3.45×104 s–1) are slow, which will lead to an enrichment in HCOO species on the surface. Figure 6(B) shows that Keq has a similar tendency to that of kf: R8 >> R7 >> R6 ≈ R5. Because of the exothermic reactions, Keq of R7 and R8 decrease with an increase in temperature, whereas for R5 and R6 (endothermic), the trend is reversed. Snapshots of IS, TS, and FS of R5–R8 are shown in Fig. 7. The calculated adsorption energies for HCOO (IS5&6), HCOH (IS7), and HCO (IS8) are –4.10, –3.32, and –2.60 eV, respectively. The corresponding adsorption configurations are η2–µ2 (O,O), η1–µ2 (C), and η2–µ2 (C,O). The molecular plane of HCOO is perpendicular to the surface. For R5, deoxidation occurred, and the C–O bond lengthened from 1.27 to 1.88 Å in TS5 to yield products of HCO and O (FS5). During the dehydrogenation of R6, the H atom is transferred away from the surface to a three-faceted site (FS6). The length of the C–H bond is 1.44 Å in TS6, which is 0.33 Å longer than that in IS6. Co-adsorbed CO2 in FS6 forms an η3–µ2 (C,O,O) configuration with two O atoms on the top site and a C atom on the bridge site. Its 10


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d(O–sCo) values are 2.07 and 2.05 Å, and the d(C–sCo) values are 2.14 and 1.96 Å. For R7, the cleaving HCO–H bond in TS7 is elongated significantly from 0.98 to 1.42 Å. For R8, because of its extremely low reaction barrier, the geometry of TS8 is similar to that of IS8. The difference is that C–H is rotated slightly from 65° to 41° in TS8. The produced CO is adsorbed at a bridge site with two distinguishing d(C–sCo) values of 2.06 and 1.78 Å (FS8). 3.1.3. COOH dissociation The COOH intermediate that is produced via the initial dissociation of HCOOH will undergo further decompositions to CO and CO2 via CO–OH and COO–H scissions, respectively. For CO formation, CO–OH scission (R9) occurs directly with an activation barrier of 0.66 eV. The reaction is strongly exothermic (–1.19 eV). For CO2 formation, the original COOH (cis) transforms to trans-COOH (R10) by OH rotation, followed by COO–H bond cleavage (R11). The corresponding barriers for R10 and R11 are 0.45 and 0.67 eV, respectively, consistent with those (0.41 eV for cis-COOH → trans-COOH and 0.78 eV for trans-COOH → CO2 + H) reported by Luo et al. from the DFT calculations of methanol steam reforming on Co(0001) surface.26 R10 is almost thermoneutral, whereas R11 is mildly exothermic (–0.54 eV). The corresponding potential energy profile and the geometries of IS, TS, and FS of R9–R11 are presented in Fig. 8. COOH decompositions to CO and CO2 are thermodynamically favored and have similar activation barriers. For R9, the breaking CO–OH bond is elongated from 1.34 to 1.74 Å in TS9. Next, dehydroxylation yields a co-adsorbed intermediate of CO and OH (FS9) on the stepped surface. CO and OH 11


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adsorb on two vicinal bridge sites through atomic C and O. The corresponding d(C–sCo) are 1.83 and 1.93 Å, and d(O–sCo) are 1.88 and 1.95 Å. During CO2 formation (R10 and R11), the OH group of COOH (IS9 & 10) is activated and rotated through TS10 to yield trans-COOH (FS10). Then, the O–H bond is stretched and pointed toward the stepped surface with a prolonged O–H length of 1.34 Å in TS11. The scission of O–H yields CO2 and H on the surface with an η3–µ2 (C,O,O) configuration for CO2 (FS11). The kf and Keq values of R9–R11 are shown in Fig. 9. The magnitude of kf is on the order of R10 > R9 ≈ R11. As a result, R9 and R11 have similar reaction rates. The activated rotation of OH in the step of the COOH configuration transformation (R10) is ~24 times faster than that of R11 at 773 K. Figure 9(B) shows that the Keq of R9 is much higher than those of R11 and R10. Using 773 K as an example, it can be seen that the corresponding values are 1.38×108, 2.36×103, and 5.74×10–1, respectively. Therefore, R9 is more thermodynamically favorable than R11. Combined with the reaction rates, it is concluded that the route for COOH direct dehydroxylation to CO (R9) occurs prior to that for COOH dehydrogenation to CO2 (R10 + R11). 3.1.4. H2 and H2O formations H2 formation from two adsorbed H species has an activation barrier of 0.60 eV, which is associated with a mild endothermicity of 0.57 eV (R12). Hydrogen will also react with OH and O species to yield H2O (R13) and OH (R14), respectively. The adsorbed OH species is produced from R2 and R9, and O species is produced from R4 and R5. OH is easier to generate than O during decomposition because of its 12


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lower formation barriers. The hydrogenation of OH to H2O has a high reaction barrier of 1.62 eV, which is strongly endothermic by 1.20 eV on the surface. H2O formation on a metal Co surface is kinetically and thermodynamically difficult. The hydrogenation of O to OH has an activation barrier of 0.98 eV with a mild exothermicity of –0.32 eV. The kf and Keq values of R12–R14 are depicted in Fig. 10. The reaction rate of R12 is higher than that of R13, and its equilibrium constant is also larger than that of R13. H2 formation is preferred over H2O formation on the Co surface. The kf and Keq values of R14 are larger than those of R13, which indicates that O hydrogenation is much more rapid and easier than OH hydrogenation. The geometries of IS, TS, and FS of R12–R14 on the surface are presented in Fig. 11. For the H2 formation, two vicinal atomic H that are located at the bridge site separately (IS12) approach the yield of the top site adsorption of the H2 species (FS12). The H–H distance changes from 2.49 Å in IS12 to 1.18 Å in TS12 and finally to 0.94 Å in FS12. Similarly, H2O formation starts from two vicinal and bridge-site adsorbed H and OH (IS13). The produced H2O (FS13) is adsorbed on the top site through its O atom with a d(O–sCo) of 2.14 Å. The length of H–OH in TS13 is 1.41 Å. The distance of the vicinal H and O in O hydrogenation is decreased from 2.68 to 1.45 Å in TS14. The produced OH is located at the bridge site with two O–sCo bonds of 1.93 Å (FS14). 3.2. Reaction pathways According to the above results, the favored initial dissociation steps of HCOOH are 13


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focused on three elementary reactions: HCOO–H scission (R1), HCO–OH scission (R2), and H–COOH scission (R3). These three reactions have activation barriers below 0.5 eV. Because of the high reaction barriers of HCOO subsequent dissociations (R5 and R6), the following decompositions are focused on COOH and HCO. Previous experimental and computational work has reported that either COOH or HCO is the reactive intermediate in formic acid decomposition, while HCOO is simply a spectator.33–35 Three potentially relevant reaction routes have been identified, and the corresponding potential energy diagram is presented in Fig. 12. The former two routes involve intermediates of COOH (P1 and P2), and the last route mediates HCO (P3). P1: HCOOH → COOH → CO P2: HCOOH → COOH → trans-COOH → CO2 P3: HCOOH → HCO → CO The rate-determining step (RDS) in P1 is COOH → CO + OH, which is associated with an activation barrier of 0.66 eV. The RDS in P2 is trans-COOH → CO2 + H, which is associated with a barrier of 0.67 eV. The difference of these two highest barriers is small at 0.01 eV, which implies that P1 and P2 are energetically compatible. The highest barrier for P3 is 0.45 eV (HCOOH → HCO + OH), which is ~0.2 eV lower than that of P1 and P2. The reaction barriers of the RDS for these three pathways increase in the order P3 < P1 ≈ P2. As a result, P3 is the most energetically preferred pathway for HCOOH decomposition on the Co-stepped surface. This is consistent with our previous result for acetic acid decomposition on the Co-stepped 14


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surface.15 In DFT calculations for propanoic acid decomposition on the Pd(111) surface (without microkinetic modeling), Lu et al. also concluded that the most likely decarbonylation pathway proceeded via acid dehydroxylation followed by α–carbon dehydrogenation.36 P1 and P2 are also feasible and likely to occur because all activation barriers in the routes are lower than 0.7 eV. The formation of COOH in P1 and P2 will compete with the formation of HCO in P3 in the initial HCOOH dissociation step because the corresponding barrier is 0.26 eV compared with 0.45 eV. Therefore, detailed microkinetic modeling needs to be carried out to determine unambiguously the favorable reaction pathways. 3.3. Microkinetic modeling The conversion of HCOOH and the relative selectivity of carbon-containing products at different temperatures as calculated by microkinetic modeling are presented in Fig. 13. HCOOH is completely converted between 300 and 1000 K. The products are unsaturated intermediates of COOH, HCOO, HCOH, HCO, and saturated species of CO and CO2. Figure 13 shows that the dominant product of HCOOH decomposition at low temperature is COOH rather than HCOO, which confirms the quite readily of H–COOH scission. This result agrees well with those reported by Lu et al.30 in their detailed microkinetic modeling study of propanoic acid decomposition. They found that the decarbonylation pathway, which starts with the α–carbon dehydrogenation of the acid, is one order of magnitude faster than the decarboxylation pathway, which 15


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begins with the O–H scission. The selectivity of HCO is kept to zero over the entire temperature range. As discussed above, the barrier for HCO dissociation to CO is extremely low (0.03 eV), and its kf is large, which indicates that HCO will be transferred immediately to CO when it is generated. However, CO in Fig. 13 is generated after 450 K. Therefore, we deduce that few to no HCO intermediates are produced during the decomposition. Thus, HCOOH decomposition is favored to undergo P1/P2 rather than P3. In addition, the HCOH selectivity is also zero, which indicates that no HCOH is produced due to its higher formation barrier (R4, Eac = 0.79 eV) over HCO formation (R2, Eac = 0.45 eV). Consequently, products of the initial dissociation of HCOOH are limited to COOH and HCOO, as shown in Fig. 13. The COOH selectivity decreases from 95% to zero at 300–1000 K whereas, in general, the HCOO selectivity increases from 5% to 42%, which suggests that HCOO–H scission becomes more competitive with H–COOH scission at high temperatures. The increasing HCOO selectivity implies the hard conversion of HCOO, which results in its surface accumulation. This agrees well with the above DFT results of HCOO formation, which has a low barrier of 0.36 eV (R1), and HCOO dissociations, which have barriers higher than 1.3 eV (R5 and R6). CO and CO2 are two products of COOH conversion. The sharp decrease of COOH selectivity occurs at 450–750 K coupled with the rapid formation of CO (R9) and the smooth formation of CO2 (R10+R11). The selectivity of CO first increases from 1% at 450 K to a maximum of 52% at 750 K and then decreases slightly to 47% at 1000 16


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K. The maximum CO2 selectivity is 12% at 750 K, which is less than 1/4 of the content of CO. Therefore, CO formation is much easier and more rapid than CO2 formation, suggesting that P1 is preferred over P2. Therefore, we conclude that the most likely pathway for HCOOH decomposition proceeds via HCOOH → COOH → CO (P1). Lu et al. also found that the most favorable pathway for propanoic acid decomposition started with the α–carbon dehydrogenation of the acid, followed by dehydroxylation.30 4. Conclusion DFT calculations were performed to determine ZPE-corrected activation barriers and reaction energies and the rate and equilibrium constants of 14 elementary reactions of formic acid decomposition over a Co-stepped surface. The DFT results show that the

dissociation likely begins with the dehydrogenation and

dehydroxylation of HCOOH to yield COOH, HCOO, and HCO, with activation barriers of 0.26, 0.36, and 0.45 eV, respectively. The corresponding formation rate constants are 1.78×1011, 1.03×1011, and 1.46×1010 s–1 at 773 K, increasing in the order of COOH > HCOO > HCO. HCOO generation is thermodynamically favored, but it is energetically difficult to convert. Conversions for HCO and COOH are readily achievable. HCO dehydrogenation to CO has an extremely low barrier of 0.03 eV. COOH conversions to CO and CO2 are two energetically similar pathways with the highest barriers of 0.66 and 0.67 eV, respectively. Microkinetic modeling shows that the dominant product of HCOOH decomposition at low temperature is COOH. With an increase in temperature, the selectivity of 17


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HCOO increases, while the selectivity of COOH decreases. The conversion of COOH to CO is preferred over CO2. By coupling the DFT and microkinetic modeling results, the most favorable formic acid decomposition pathway is identified as HCOOH → COOH → CO. Acknowledgments The authors acknowledge financial support from the National Natural Science Foundation of China (Nos. 51406090 and 51406093), the Ningbo Natural Science Foundation of China (No. 2014A610119), and the K.C. Wong Magna Fund in Ningbo University. We are also grateful for the Special Program for Applied Research on Super Computation of the NSFC-Guangdong Joint Fund (the second phase). References (1) Wang, D.; Czernik, S.; Montané, D.; Mann, M.; Chornet, E. Ind. Eng. Chem. Res 1997, 36, 1507-1518. (2) Wang, S. R.; Zhang, F.; Cai, Q. J.; Li, X. B.; Zhu, L. J.; Wang, Q.; Luo, Z. Y. Int. J. Hydrogen Energy 2014, 39, 2018-2025. (3) Bridgwater, A. V. Biomass Bioenergy 2012, 38, 68-94. (4) Wang, D.; Czernik, S.; Chornet, E. Energy Fuels 1998, 12, 19-24. (5) Rioche, C.; Kulkarni, S.; Meunier, F. C.; Breen, J. P.; Burch, R. Appl. Catal., B 2005, 61, 130-139. (6) Trane, R.; Dahl, S.; Skiøth-Rasmussen, M. S.; Jensen, A. D. Int. J. Hydrogen Energy 2012, 37, 6447-6472. (7) Wang, S. R.; Li, X. B.; Zhang, F.; Cai, Q. J.; Wang, Y. R.; Luo, Z. Y. Int. J. 18


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Hydrogen Energy 2013, 38, 16038-16047. (8) Xing, R.; Dagle, V. L.; Flake, Mathew.; Kovarik, L.; Albrecht, K. O.; Deshmane, Chinmay.; Dagle, R. A. Catal. Today 2016, 269, 166-174. (9) Remiro, A.; Valle, B.; Aguayo, A. T.; Bilbao, J.; Gayubo, A. G.; Fuel Process. Technol. 2013, 115, 222-232. (10) Mohan, D.; Pittman, C. U.; Steele, P. H. Energy Fuels 2006, 20, 848-889. (11) Guo, Z. G.; Wang, S. R.; Gu, Y. L.; Xu, G. H.; Li, X.; Luo, Z. Y. Sep. Purif. Technol. 2010, 76, 52-57. (12) Wang, S. R.; Li, X. B.; Guo, L.; Luo, Z. Y. Int. J. Hydrogen Energy 2012, 37, 11122-11131. (13) Galdámez, J. R.; García, L.; Bilbao, R. Energy Fuels 2005, 19, 1133-1142. (14) Wang, S. R.; Gu, Y. L.; Liu, Q.; Yao, Y.; Guo, Z. G.; Luo, Z. Y.; Cen, K. F. Fuel Process. Technol. 2009, 90, 738-745. (15) Li, X. B.; Wang, S. R.; Zhu, Y. Y.; Yang, G. H.; Zheng, P. J. Int. J. Hydrogen Energy 2015, 40, 330-339. (16) Wang, Q.; Wang, S. R.; Li, X. B.; Guo, L. BioResources 2013, 8, 2897-2909. (17) Hu, X.; Lu, G. X. Chem. Lett. 2006, 35, 452-453. (18) Yoo, J. S.; Abild-Pedersen, F.; Nørskov, J. K.; Studt, F. ACS Catal. 2014, 4, 1226-1233. (19) Luo, Q. Q.; Feng, G.; Beller, Matthias.; Jiao, H. J. J. Phys. Chem. C 2012, 116, 4149-4156. (20) Singh, S.; Li, S.; Carrasquillo-Flores, R.; Alba-Rubio A. C.; Dumesic, J. A.; 19


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Mavrikakis, M. AIChE J. 2014, 60, 1303-1319. (21) Hu, C. Q.; Ting, S. W.; Chan, K. Y.; Huang, W. Int. J. Hydrogen Energy 2012, 37, 15956-15965. (22) Honkala, K.; Hellman, A.; Remediakis I. N.; Logadottir, A.; Carlsson, A.; Dahl, S.; Christensen, C. H.; Nørskov, J. K. Science 2005, 307, 555-558. (23) Kresse, G.; Hafner, J. Phys. Rev. B 1993, 47, 558-561. (24) Kresse, G.; Furthmuller, J. J. Comp. Mater. Sci. 1996, 6, 15-50. (25) Kresse, G.; Furthmuller, J. Phys. Rev. B 1996, 54, 11169-11186. (26) Luo, W. J.; Asthagiri, A. J. Phys. Chem. C 2014, 118, 15274-15285. (27) de la Peña O’Shea, V. A.; de P. R. Moreira, I.; Roldán, A.; Illas, F. J. Chem. Phys. 2010, 133, 024701. (28) Henkelman, G.; Uberuaga, B. P.; Jonsson, H. J. Chem. Phys. 2000, 113, 9901-9904. (29) Henkelman, G.; Jonsson, H. J. Chem. Phys. 2000, 113, 9978-9985. (30) Lu, J. M.; Behtash, S.; Faheem, M.; Heyden, A. J. Catal. 2013, 305, 56-66. (31) Scaranto, J.; Mavrikakis, M. Surf. Sci. 2016, 648, 201-211. (32) Scaranto, J.; Mavrikakis, M. Surf. Sci. 2016, 650, 111-120. (33) Wilhelm, S.; Vielstich, W.; Buschmann, H.; Iwasita, T. J. Electroanal. Chem. 1987, 229, 377-384. (34) Neurock, M.; Janik, M.; Wieckowski, A. Faraday Discuss. 2008, 140, 363-378. (35) Chen, Y. X.; Heinen, M.; Jusys, Z.; Behm, R. J. Langmuir 2006, 22, 10399-10408. 20


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(36) Lu, J. M.; Behtash, S.; Heyden, A. J. Phys. Chem. C 2012, 116, 14328-14341.

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List of Tables: Table 1. ZPE-corrected activation energies (Eac), reaction energies (∆H), and rate parameters for all elementary reaction steps in formic acid decomposition over a Co-stepped surface.

List of Figures: Fig. 1. Scheme of possible HCOOH decomposition pathways. Fig. 2. Potential energy profile of initial HCOOH dissociation. Fig. 3. Temperature dependence of the forward rate constant and equilibrium constant of initial HCOOH dissociation. Fig. 4. Side and top views of IS, TS, and FS of steps R1−R4 on a Co-stepped surface. Blue, brown, red, and white spheres represent Co, C, O, and H, respectively. Fig. 5. Potential energy profile of HCOO and HCOH dissociations. Fig. 6. Temperature dependence of the forward rate constant and equilibrium constant of HCOO and HCOH decompositions. Fig. 7. Side and top views of IS, TS, and FS of steps R5−R8. Fig. 8. Potential energy profile of COOH dissociation. Fig. 9. Temperature dependence of the forward rate constant and equilibrium constant of COOH dissociation. Fig. 10. Temperature dependence of the forward rate constant and equilibrium constant of H2, H2O, and OH formations. Fig. 11. Side and top views of IS, TS, and FS of steps R12−R14. 22


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Fig. 12. Potential energy surface of favorable HCOOH decomposition pathways. Fig. 13. Conversion and product selectivity of HCOOH decomposition on a Co-stepped surface using microkinetic modeling.

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Table 1. ZPE-corrected activation energies (Eac), reaction energies (∆H), and rate parameters for all elementary reaction steps in formic acid decomposition over a Co-stepped surface. T = 773 K Steps

Reactions

Eac (eV)

∆H (eV) kf (s–1)

Keq

R1

HCOOH → HCOO + H

0.36

–1.12

1.03×1011

1.27×107

R2

HCOOH → HCO + OH

0.45

–0.62

1.46×1010

2.19×104

R3

HCOOH → COOH + H

0.26

–0.42

1.78×1011

6.33×102

R4

HCOOH → HCOH + O

0.79

–0.19

5.96×107

2.06×101

R5

HCOO → HCO + O

1.30

0.66

3.45×104

9.91×10–5

R6

HCOO → CO2 + H

1.33

0.59

3.97×104

3.84×10–4

R7

HCOH → HCO + H

0.61

–0.42

1.11×109

4.99×102

R8

HCO → CO + H

0.03 a

–0.97

4.99×1012

1.62×106

R9

COOH → CO + OH

0.66

–1.19

1.20×109

1.38×108

R10

COOH → trans-COOH

0.45

0.04

1.38×1010

5.74×10–1

R11

trans-COOH → CO2 + H

0.67

–0.54

5.82×108

2.36×103

R12

H + H → H2

0.60

0.57

4.03×109

4.52×10–4

R13

H + OH → H2O

1.62

1.20

4.88×102

6.83×10–8

R14

O + H → OH

0.98

–0.32

7.61×106

1.34×102

a

Activation energy with no ZPE correction.

24


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Fig. 1. Scheme of possible HCOOH decomposition pathways.

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Fig. 2. Potential energy profile of initial HCOOH dissociation.

26


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Fig. 3. Temperature dependence of the forward rate constant and equilibrium constant of initial HCOOH dissociation.

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IS

FS1

TS1

FS2

TS2

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FS3

TS3

FS4

TS4

Fig. 4. Side and top views of IS, TS, and FS of steps R1−R4 on a Co-stepped surface. Blue, brown, red, and white spheres represent Co, C, O, and H, respectively.

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Fig. 5. Potential energy profile of HCOO and HCOH dissociations.

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Fig. 6. Temperature dependence of the forward rate constant and equilibrium constant of HCOO and HCOH decompositions.

30


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IS5&6

IS7

FS8

IS8

TS5

FS5

TS6

FS6

TS7

FS7

TS8

Fig. 7. Side and top views of IS, TS, and FS of steps R5−R8.

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Fig. 8. Potential energy profile of COOH dissociation.

32


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Fig. 9. Temperature dependence of the forward rate constant and equilibrium constant of COOH dissociation.

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Fig. 10. Temperature dependence of the forward rate constant and equilibrium constant of H2, H2O, and OH formations.

34


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IS12

TS12

FS12

IS14

IS13

TS14

TS13

FS14

Fig. 11. Side and top views of IS, TS, and FS of steps R12−R14.

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FS13

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Fig. 12. Potential energy surface of favorable HCOOH decomposition pathways.

36


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Fig. 13. Conversion and product selectivity of HCOOH decomposition on a Co-stepped surface using microkinetic modeling.

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Density Functional Theory and Microkinetic Studies of Bio-oil

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