Research Article pubs.acs.org/journal/ascecg
Cite This: ACS Sustainable Chem. Eng. 2017, 5, 10371-10378
Quantum Chemistry Analysis of Reaction Thermodynamics for Hydrogenation and Hydrogenolysis of Aromatic Biomass Model Compounds Laurene Petitjean,‡ Raphael Gagne,† Evan S. Beach,‡ Jason An,† Paul T. Anastas,*,‡ and Dequan Xiao*,† †
Center for Integrative Materials Discovery, Department of Chemistry and Chemical Engineering, University of New Haven, 300 Boston Post Road, West Haven, Connecticut 06516, United States ‡ Center for Green Chemistry and Green Engineering, Yale University, 370 Prospect Street, New Haven, Connecticut 06520, United States S Supporting Information *
ABSTRACT: Designing effective and selective reactions at sustainable or mild conditions is key for the valorization or refinery of lignin biomass using H2 reduction methods. However, it remains unclear what are the feasible mildest conditions for the reductive valorization of lignin, at which transformations can be designed. Here, we aim to exploit this critically important question using quantum chemistry calculations to systematically analyze the thermodynamics of hydrogenation and hydrogenolysis of typical functional groups found in lignin based on a set of aromatic model compounds. Our results show that it is thermodynamically feasible to break ether linkages and remove oxygen content in the model compounds even at room temperature, room pressure, and in aqueous solvent (i.e., the global mildest conditions). Interestingly, the potential influence on the thermodynamics by reaction variables is ranked in the order of temperature > H2 pressure > solvent dielectric constant; a strategically chosen solvent may enable increased selectivity for hydrogenolysis over hydrogenation. Our predicted reaction thermodynamics is consistent with our experimental findings of probed reaction pathways. This work may inspire researchers to pursue the design of “ultimate” green biomass conversion processes closer to the global mildest conditions. KEYWORDS: Gibbs free energy, Lignin, Hydrogenolysis, Hydrogenation, Biomass
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INTRODUCTION As a renewable carbon source, biomass (consisting of mainly lignin, cellulose, and hemicellose) has the potential to provide energy and platform chemicals,1 therefore mitigating society’s unsustainable consumption of finite resources, and decreasing the environmental impacts of the latter’s extraction, distribution, and use.2,3 Lignin comprises up to 25−35% of plant biomass and consists of methoxylated and hydroxylated 4propylphenol units coupled as ethers and occasionally crosslinked with carbon−carbon bonds.4 Whereas cellulose and hemicellulose have been extensively studied for the generation of bioethanol and other platform chemicals, lignin remains a challenge to valorize and is currently still discarded or burnt after separation from the glycopolymers. The valorization of lignin, however, could provide a renewable source of aromatic monomers, which could serve as drop-in replacements to those that are petroleum derived. Designing processes that work under sustainable conditions is the key for the valorization or refinery of lignin using H2 reduction methods.5−7 Reductive approaches to depolymerize lignin are advantageous for their © 2017 American Chemical Society
ability to quench char-forming radical species, limit unwanted repolymerization, and thus potentially increase the yield and selectivity of targeted aromatic products. Efficient and selective approaches reduce economic and environmental costs associated with the production of waste and purification or separation processes. The splitting of water, furthermore, can allow hydrogen to be renewably sourced. However, it remains unclear what are the feasible mild conditions for the reductive valorization of lignin for which methodologies can be designed. Here, we aim to exploit this critically important question using quantum chemistry calculations to systematically analyze the reaction thermodynamics of hydrogenation and hydrogenolysis based on a set of aromatic model compounds representing lignin. Our goal is to evidence the mildest thermodynamic conditions (lowest energy input, most benign solvent, least Received: July 18, 2017 Revised: September 15, 2017 Published: October 9, 2017 10371
DOI: 10.1021/acssuschemeng.7b02384 ACS Sustainable Chem. Eng. 2017, 5, 10371−10378
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ACS Sustainable Chemistry & Engineering
study reactions of lignin.10 These compounds are also found as reactive lignin depolymerization intermediates and are thus important to study. We present the two-dimensional (2D) contours plots of Gibbs free energy (ΔrG) of reaction as a function of varying PH2 and T, εs and T, and εs and PH2, respectively. We show these functional groups have different thermodynamic properties in terms of reactions with hydrogen, and our calculated thermodynamics results are consistent with the experimental studies of the model compounds’ H2reductive reactions with or without a Cu-doped porous metal oxide (Cu-PMO) catalyst. Cu-PMO has the advantage of being a heterogeneous catalyst composed of earth abundant elements. It is thus easily recyclable, easier to implement on industrial scale, and relatively cheap to synthesize. This work provides the theoretical foundation for the feasibility of designing more sustainable processes for lignin valorization, potentially utilizing green catalysts that function nearer the global mildest conditions. We assume that introducing catalysts to the reactions will not change reaction thermodynamics, but affect the reaction kinetics, leading to potential expedition of reactions. In addition, we will identify possible reaction pathways using known catalysts (Cu-PMO) in experiments, serving as evidence of the thermodynamic feasibility of specific reactions. This study does not focus on exploring the detailed kinetics of catalysis and the design of actual catalysts, which will be the subject of future work.
reagent inputs) at which these transformations can be effected for lignin valorization. Reaction of lignin with hydrogen can produce small organic molecules through the breakage of aromatic ether bonds (dominant) or C−C bonds, lower the oxygen content, and increase the hydrogen/carbon ratio, a valuable route for the production of biofuels. Three key reaction variables: temperature (T), H2 pressure (PH2), and solvent dielectric constants (εs) play a significant role in reaction efficiency and selectivity and are commonly and easily altered to optimize reactions. However, there remains a lack of knowledge of the thermodynamics of hydrogenation or hydrogenolysis of lignin with systematic variation of T, PH2, and εs. The early literature has some interesting findings for the thermodynamics of hydrogenation of benzene8 that is relevant to lignin model compounds. For example, the formation of cyclohexa-1,3-diene by adding one H2 molecule to benzene has a positive Gibbs free energy in the temperature range of 298.15−1003 K, while the formation of cyclohexane by adding three H2 molecules has a negative Gibbs free energy.8 In this work, we used quantum chemistry calculations at the density functional theory (DFT) level9 to study reaction thermodynamics for hydrogenation or hydrogenolysis of six typical functional groups in lignin: benzylic aldehyde (F1), aliphatic double-bond (F2), aromatic double-bond (F3), benzylic ketone (F4), aromatic ether (F5), and benzylic hydroxyl (F6) (Scheme 1a), with the systematic variation of T, PH2, and εs. These functional groups were represented in four biomass model compounds for the calculations (Scheme 1b). The complexity and variability of lignin polymers has stimulated the use of simple lignin model compounds to
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EXPERIMENTAL SECTION
The high-performance computational chemistry software NWChem11 was used to determine the energy of different reaction pathways of the hydrogenation or hydrogenolysis of four biomass model compounds. The structures were built in .xyz format using the model building program Avogadro.12 The .xyz geometries were then used in NWChem to obtain their respective optimized geometries, as seen in section 4 of the Supporting Information (i.e., SI). This ensured the intermolecular energy minimization of each of the structures. The optimized structures were then used to perform vibrational frequency analysis and solvation energy calculations. All DFT calculations were performed at the B3LYP/6-31G* level. We performed additional calculations with different basis sets (6-311g* and cc-pvdz) for the aldehyde moiety of the vanillin model compound (F1), at 1 bar and 571.15 K (see SI Table S2) to confirm our method. Gibbs energy calculations in continuum solvation model showed an error of 5 kcal/ mol for neutral organic compounds.13 This general error is smaller than all results discussed in the present paper, unless otherwise specified. The output files of the frequency calculations provided the following information: zero-point correction to energy (E1), thermal correction to enthalpy (H), and total entropy (S). The solvation calculation provided the total DFT energy (E0) and the electrostatic solvation energy (Es). Equation 1 was used to determine the change in Gibbs free energy (G).
Scheme 1. (a) Six Functional Groups for the Quantum Chemistry Thermodynamics Study of Hydrogenation (F1− F3) and Hydrogenolysis (F4−F6), where R or R’ Denotes Aliphatic or Aromatic Functional Groups in Lignin Polymers or Model Compounds and (b) Four Lignin Model Compounds Used to Study the Six Functional Groupsa
(1)
dG = H − TS
This information was then used to find Gibbs free energy (G) of each structure determined by eq 2. G = dG + E0 + E1 + Es
(2)
The Gibbs free energy of each structure was then used to determine the total change in Gibbs free energy for each reaction pathway. For example, the hydrogenation reaction can be generalized by the following chemical equation in solution
R(solu) + H 2(solu) ↔ RH 2(solu) a
(3)
The Gibbs free energy change of this reaction in gas phase R(g) + H2(g) ↔ RH2(g) is
In part b, the waved line denotes the position of bond cleavage. 10372
DOI: 10.1021/acssuschemeng.7b02384 ACS Sustainable Chem. Eng. 2017, 5, 10371−10378
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ACS Sustainable Chemistry & Engineering ΔGg = Gg,RH2 − Gg,H2 − Gg,R
based on 2 (vanillyl alcohol); F2 (aliphatic double-bond) is based on 3 (eugenol); and F4 (ketone) is based on 4 (acetovanillone). We studied the thermodynamics of the hydrogenation reactions of F1−F3 and the hydrogenolysis reactions of F4−F6, respectively (Scheme S1, c.f. the SI). We calculated ΔrG for the hydrogenation and hydrogenolysis reactions of F1−F6 in the ranges of 1 ≤ PH2 ≤ 60 bar, 273.15 ≤
(4)
After that, reaction ΔrG in solution is computed using the following Born−Haber cycle:The solvation energy ΔGs,i for each species is
T ≤ 573.15 K, and 0 ≤ εs ≤ 80 which, in our view, represents the most practical and least energy intensive conditions. computed using the COSMO solvation model.14 The reaction free energy in solution ΔGsolu is computed by ΔGsolu = ΔGg + ΔGs,RH2 − ΔGs,R − ΔGs,H2
(6)
Here, for the simplicity of calculation, we assume that varying the partial pressure of H2 does not influence the activities of R and RH2 (biomass model compounds before and after hydrogenation). We allow the partial pressure of H2 to vary in the range of 1−60 bar. In eq 4, the thermally corrected Gibbs energy of H2 is pH ° 2 + RT ln 2 Gg,H2(p) = Gg,H p° (7) In eq 6, we also assume that the solvation energies (per mole) of the reactants are not influenced by PH2, even for H2 itself. As shown in the literature,15 in a very dilute solution of a gas (e.g., H2 here), the solvation energy of the gas molecule is related to Henry’s constant only, which does not change with PH2. In eq 6, the temperature T influences the solvation energies. To first study this effect of temperature, we computed the dielectric constant εs of the solvent at a particular temperature T. We found that εs of methanol decays exponentially with the increase of temperature (see Figure S1 in the SI). Hence, we used methanol as a model solvent to study the influence of T on ΔrG. For the sake of generalization, we continued the trend of exponential decay for εs beyond the temperature of 300 K. Once we obtained εs (at T) based on the fitting experimental trend, we then used that εs to compute the solvation energy at the particular temperature T using the COSMO solvation model. Detailed descriptions of experimental procedures can be found in the SI. Cu-PMO is synthesized and characterized as described in the literature.10 Copper constitutes 20 mol % of M2+, with M2+:M3+ kept at 3:1. Elemental analyses showed that the metals were incorporated in the anticipated amounts, furnishing a catalyst with metal ratios of Cu0.57Mg2.25Al1.00. XRPD measurements indicate that Cu-PMO loses the hydrotalcite-like structure, becoming an amorphous material, after calcination in air for 24 h at 733.15 K. Cu-PMO is known to have a surface area of ∼137 m2/g through BET analyses. Biomass model compounds 1−4 (see Scheme 1b) and solvent (methanol) were purchased from EMD Millipore, Sigma-Aldrich, AlfaAesar, Merck, or TCI and used as received. All hydrogenation reactions were set up in a 100 mL stainless-steel Parr reactor equipped with a mechanical stirrer. The reactions were pressurized under hydrogen atmosphere (Tech Air, Ultra High Purity). The loaded reactor was placed on a benchtop Parr stand equipped with a Parr 4843 reactor controller.
Figure 1. Reaction Gibbs free energies changes versus H2 pressure and temperature for the hydrogenation and hydrogenolysis of six functional groups. Contours a−f correspond to the reaction Gibbs free energies of functional groups F1−F6, respectively.
Figure 1 demonstrates the effect of varying T and PH2. At PH2 = 1 bar, hydrogenation of F1 (benzylic aldehyde) and F3 (aromatic double-bond) have a positive ΔrG, while the other reactions have a negative ΔrG. Figure 1 calculations are perfomed using methanol as representative solvent with a dielectric constant value corrected for its change with temperature (please refer to Figure S1 in the SI). The predicted positive ΔG of adding one H2 to the phenyl ring is consistent with the experimental finding for benzene hydrogenation.8 However, the reduction of benzene ring to with the addition of 2 or 3 H2 shows a negative ΔG, and thus is thermodynamically possible.8 This is consistent with the literature that the aromatic ring hydrogenation of lignin model compounds, and of products from the depolymerization of lignin itself, have been reported using heterogeneous nickel catalysts at T and PH2 within the range considered herein.16 We envision that, in accordance with calculations on benzene hydrogenation here, reduction of the cyclohexadiene derivative, resulting from addition of one hydrogen equivalent, to cyclohexene or cyclohexane is thermodynamically favored
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RESULTS AND DISCUSSION We analyzed the thermodynamics of hydrogenation or hydrogenolysis for six typical functional groups in lignin polymers or derivatives (Scheme 1a), integrated in biomass model compounds 1−4 shown in Scheme 1b. F1 (benzylic aldehyde), F3 (aromatic double-bond), and F5 (aromatic ether) are based on 1 (vanillin); F6 (benzylic hydroxyl) is 10373
DOI: 10.1021/acssuschemeng.7b02384 ACS Sustainable Chem. Eng. 2017, 5, 10371−10378
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ACS Sustainable Chemistry & Engineering relative to the aromatic substrate, especially due to the release of the ring strain arising from loss of aromaticity. Overall, varying PH2 or T does not change the sign of ΔrG for any of six functional groups. Thus, in the ranges of PH2 (1−60 bar) and T (273.15−573.15 K), the hydrogenation of F1 (benzylic aldehyde) and F3 (the first aromatic double-bond) are not thermodynamically favored; and the hydrogenolysis of F2 (aliphatic double bond), F4 (ketone), F5 (aromatic ether), and F6 (benzylic hydroxyl) are spontaneous. With the increase of T (at constant PH2), ΔrG increases for the hydrogenation reactions of F1−F3 (Figure 1) because these hydrogenation reactions involve a negative change in entropy (for F1, ΔS = −17.96 cal/mol·K, for F2, ΔS = −27.67 cal/mol· K; whereas for F3, ΔS = −20.55 cal/mol·K). Consequently, ΔrG becomes directly proportional to T (rather than to negative T) and will increase with temperature. In other words, increasing T reduces the extent of reduction. In contrast, for the hydrogenolysis reactions of F5−F6, ΔrG decreases with the increase of T (at constant PH2), as the hydrogenolysis reactions are endothermic. Thus, increasing T accelerates the extent of hydrogenolysis for aromatic ether and hydroxyl (aliphatic or aromatic) groups. The hydrogenolysis of F4 (ketone) is a combined reaction of hydrogenation of carbonyl group and hydrogenolysis of benzylic hydroxyl group, and the overall ΔrG is negative, due to the large negative ΔrG from the hydrogenolysis of benzylic hydroxyl group. This effect can be attributed to a large entropy gain from the release of a water molecule from hydrogenolysis of the benzyl alcohol. Specifically, the entropy of F4 is ΔS = +18.24 cal/mol·K, thus leading to a large negative ΔrG. Increasing temperature, however, leads to the increase of ΔrG for F4, as the hydrogenation of the carbonyl group dominates its temperature-dependent behavior. With the increase of PH2 (at constant T), ΔrG for each reaction decreases. As expected from Le Chatelier’s principle,17 increasing PH2 favors both hydrogenation and hydrogenolysis reactions. Figure 1 also shows, as expected, that T has a higher influence to ΔrG than PH2. For example, in Figure 1b, with a 2fold increase in PH2, the change of ΔrG is less than 1 kcal/mol. In contrast, with a 2-fold increase in T, the change of ΔrG is about 3−4 kcal/mol. Figure 2 shows the variations of ΔrG with the change of εs (0−80) and T (273.15−573.15 K). The dielectric constant in Figure 2 represents different solvents and was calculated as independent from temperature. It is our aim that a solvent with a particular dielectric constant at a particular temperature can be chosen by experimentalists to target a thermodynamic window where the desired reaction occurs. At constant εs, varying T causes a change of 4−9 kcal/mol in ΔrG for all of the reactions. We notice that εs causes only subtle changes in ΔrG, compared to the influence of T. First, variation of εs (0−80) does not change the sign of ΔrG in any of the reactions except for the hydrogenolysis of F3 (ketone). For the hydrogenolysis of F4, ΔrG becomes slightly positive when εs < 5 and T > 473.15 K. This could be attributed to the decreased solubility of the released water molecule from hydrogenolysis of F4 in solvents with low dielectric constants. Second, for the hydrogenation reactions of F1−F3, increasing εs from 0 to 80 causes a change of 1 kcal/mol in ΔrG. Meanwhile, for the hydrogenolysis reactions of F5−F6, εs does not cause a
Figure 2. Reaction Gibbs free energies changes versus dielectric constant and temperature for the hydrogenation and hydrogenolysis of six functional groups, where H2 pressure is 1 bar. Contours a−f correspond to the reaction Gibbs free energies of functional groups F1−F6, respectively.
variation of more than 1 kcal/mol in ΔrG in a wide range of εs (5 < εs < 80), which is not large enough to be significant. Analogous to the F4 reaction, the solubility of catechol (from F5) and water (from F6) yielded from hydrogenolysis will be limited in solvents with very low dielectric constants, thus increasing the change in Gibbs free energy within the range εs < 5. As shown by all the ΔrG contours here, it is clear that εs plays solely a subtle role in influencing reaction thermodynamics. This is a helpful insight for selection of green solvents for biomass conversion. For example, it suggests further efforts should be made to use aqueous solvent systems, building on recent work on hydrotreating of biomass molecules.18 For all the reactions, the optimal εs (where ΔrG is the lowest) are different. For example, as shown in Figure 2, the most favored εs is 0 for F1, 0 for F2, 5 for F3, 80 for F4, 38 for F5, and 80 for F6. The corresponding solvents with suitable εs values can be found in the literature.19 For example, at 293.15 K, n-butyl acetate has εs of 5, formic acid has εs of 58, and water has εs of 80. It is interesting that ΔrG increases by 2 kcal/mol when εs < 5 for F4−F6. Thus, the hydrogenation reactions (for F1−F3) favors low εs, while the hydrogenolysis reactions (for F4−F6) favored high εs. This can be understood from the polarity change of chemical species in the reaction. The difference in solvation energies between reactants and products for each reaction (F1−F6), in methanol at 303 K, are expressed in the experimental Supporting Information (Table S3). Results show solely slight changes ( PH2 > εs. Even though εs could be important in influencing reaction kinetics and selectivity by biasing the solvation of chemical species, it plays only a subtle role in changing ΔrG for the reactions here when εs > 5. This is a helpful insight for selection of green solvents for biomass conversion. In the ranges of T (273.15−573.15 K), PH2 (1−60 bar), and εs (0−80), the thermodynamically favored reactions include hydrogenation of aliphatic double bonds and hydrogenolysis of carbonyl, aromatic ether, and hydroxyl groups in the set of studied aromatic model compounds. However, hydrogenation of the first aromatic double-bond in a phenyl ring is
while achieving efficiency and selectivity toward aromatic phenolics.10 The Cu-PMO catalyst can also achieve selective hydrogenation of aliphatic double-bonds and hydrogenolysis of ketones.10 The observations made with this catalyst are consistent with our prediction of the thermodynamics at 453.15 K and 40 bar PH2 (Scheme 2). In fact, hydrogenolysis of a benzylic hydroxyl (F6) which is calculated to be thermodynamically favored but had not been detected in experiments without catalysts, provides product in the presence of Cu-PMO, evidencing that a large kinetic barrier was limiting our ability to observe this considered reaction (Scheme 2). We do anticipate that improved catalyst design could enable the predicted reaction pathway. As one of the applications, our predicted thermodynamics of hydrogenation or hydrogenolysis can provide new insights for the design of processes for lignin degradation. For example, Figure 4 illustrates the experimental working conditions of PH2 and T for lignin depolymerization for a set of known metallic catalysts, including heterogeneous noble metals, heterogeneous transition metals and homogeneous metallic catalysts. Near the global mildest condition (i.e., room temperature and room pressure), we define an “ideal” region where PH2 = 1−10 bar and T = 298.15−393.15 K. So far, no process utilizing heterogeneous catalysts fall in the “ideal” region. Hence, it is a 10376
DOI: 10.1021/acssuschemeng.7b02384 ACS Sustainable Chem. Eng. 2017, 5, 10371−10378
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ACS Sustainable Chemistry & Engineering Scheme 3. Alternative Pathway for the Formation of Vanillyl Alcohol from Vanillin
thermodynamically disfavored in the studied model compounds. Based on the feasibility of hydrogenolysis of aromatic ether in vanillin, we anticipate that it may be feasible to perform lignin depolymerization and upgrading even at the global mildest condition through the breaking of aromatic ether linkages and the removal of oxygen contents via hydrogenolysis of carbonyl and hydroxyl groups. Currently known methods for lignin depolymerization do not yet reach the predicted mildest conditions. Our study provides the theoretical foundation for the feasibility of designing novel reactions, for which new green catalysts (e.g., transition metal catalysts) will be crucial, for lignin conversion near the global mildest conditions, i.e., room temperature, room pressure, and aqueous solvent. Our work will hopefully inspire and enable the design and development of new-generation processes for biomass valorization.
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tions, 1H NMR spectra of crude reduction reaction mixtures, optimized geometries for model compounds, characterization of isolated reaction products, temperature dependence of the dielectric constant for methanol, reaction Gibbs Free Energy for the hydrogenation of vanillin with varying basis sets, and solvation free energies for reactions F1−F6 in methanol at 303 K (PDF)
AUTHOR INFORMATION
Corresponding Authors
*E-mail:
[email protected] (D.X.). *E-mail:
[email protected] (P.T.A.). ORCID
Laurene Petitjean: 0000-0001-7730-6655 Paul T. Anastas: 0000-0003-4777-5172 Funding
The summer research grant and research fellowship (2014 and 2015) of the University of New Haven, the faculty start-up fund of the University of New Haven, and QAFCO helped fund this work. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS D.X. thanks the faculty start-up fund and summer research grant and research fellowship (in 2014 and 2015) of the University of New Haven. P.T.A. would like to thank QAFCO for continued funding support. The Center for Green Chemistry and Green Engineering is thankful to the School of Forestry and Environmental Studies for its support.
Figure 4. Working conditions of a number of known catalysts in the literature: Cu/Cr,21 Ni/Mo/S,22 Cu/Fe/Sn/S,23 Cu-PMO 1,24 Pd/ C,25 PdC+ZnCl2,26 and Cu-PMO 2.27
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ASSOCIATED CONTENT
S Supporting Information *
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The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acssuschemeng.7b02384. General experimental details, synthesis of Cu-PMO catalyst, synthesis of 2-methoxy-4-(methoxymethyl)phenol, general procedure for reduction of lignin model compounds, model compound reduction reac-
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