Impact of pH on Aqueous-Phase Phenol Hydrogenation Catalyzed by

Dec 19, 2018 - In aqueous phase, the rates of hydrogenation of aromatic substrates such as phenol on Pt/C and Rh/C are influenced by varying activity ...
0 downloads 0 Views 400KB Size
Download PDF
Subscriber access provided by University of Winnipeg Library

Article

Impact of pH on Aqueous-Phase Phenol Hydrogenation Catalyzed by Carbon-Supported Pt and Rh Nirala Singh, Mal Soon Lee, Sneha A Akhade, Guanhua Cheng, Donald M. Camaioni, Oliver Y. Gutiérrez, Vassiliki-Alexandra Glezakou, Roger Rousseau, Johannes A. Lercher, and Charles T. Campbell ACS Catal., Just Accepted Manuscript • DOI: 10.1021/acscatal.8b04039 • Publication Date (Web): 19 Dec 2018 Downloaded from http://pubs.acs.org on December 19, 2018

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 9 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

Impact of pH on Aqueous-Phase Phenol Hydrogenation Catalyzed by Carbon-Supported Pt and Rh Nirala Singh,†,‡ Mal-Soon Lee,‡ Sneha A. Akhade,‡ Guanhua Cheng,║ Donald M. Camaioni,‡ Oliver Y. Gutiérrez,‡ Vassiliki-Alexandra Glezakou,‡ Roger Rousseau,*,‡ Johannes A. Lercher*,‡,║ Charles T. Campbell*,† †Department of Chemistry, University of Washington, Seattle, Washington 98105-1700, United States ‡Institute for Integrated Catalysis, Pacific Northwest National Laboratory, Richland, Washington 99354, United States ║Department of Chemistry and Catalysis Research Center, Technische Universität München, D-85748 Garching, Germany ABSTRACT: In aqueous phase, the rates of hydrogenation of aromatic substrates such as phenol on Pt/C and Rh/C are influenced by varying activity of hydronium ions. Decreasing the pH from 8 to 1 increases the rate of hydrogenation of phenol on Pt at 20 bar H2 by 15-fold. This increase is attributed to weakening of the hydrogen binding energy (HBE) on the metal surface with decreasing pH. A weaker HBE at lower pH is also predicted by ab initio molecular dynamics simulations, providing atomistic insight into the impact of electrolyte ion distribution and interfacial solvent reorganization on HBE. The lower HBE results in a decrease in the activation energy for addition of adsorbed H from the metal to the adsorbed organic (with a Brønsted-Evans-Polanyi slope of ~1). The kinetic model derived accounts also for the lack of pH dependence at low hydrogen coverages (at 1 bar H2 on Pt or up to 70 bar H2 on Rh), when the weakening of the HBE decreases the hydrogen coverage.

KEYWORDS Hydrogenation, phenol, platinum, Brønsted-Evans-Polanyi, aqueous-phase, hydrogen adsorption, pH effects, molecular dynamics simulations

INTRODUCTION Understanding and controlling the aqueous-phase catalytic hydrogenation of organic compounds on metals is crucial to develop the optimum reaction conditions for a large number of reactions, including the reductive conversion of biomass. Unlike gas-phase hydrogenation, the impact of water, dissolved ionic reactants (also ions) and the solution/solid interface (double layer) are only beginning to be explored.1–4 Among these parameters, the impact of the thermodynamic activity of hydronium ions on the catalytic activity for hydrogenation of organic molecules is of particular importance, being linked to key reactions in the potential utilization of renewable electricity. Hydrogen evolution and oxidation (HER and HOR) are two orders of magnitude faster in acid than in base medium on platinum group metals.5–9 Various studies show non-Nernstian shifts in cyclic voltammogram peaks associated with hydrogen adsorption/desorption processes with pH,6,7,9–18 and this effect has been interpreted by some as a stronger hydrogen binding energy (HBE) or apparent HBE at higher pH for Pt group metals.6,7,9,15–17,19,20 (HBE refers to the strength of binding of H adatoms, i.e. Had, to the surface.) Yan and co-workers have suggested that the sole descriptor for pH-dependent HER/HOR kinetics follows the Sabatier principle and is primarily characterized by the strength of H binding on the metal electrode.15,16,21 We refer the readers to these recent review, research and perspective articles for a detailed

discussion on the role of pH on adsorbed H, its HBE and the catalytic HER rate.6,12,14,16,21–28 It should be noted that metals with weaker adsorption of hydrogen (free energy near zero, HBE ~−23 kJ mol-1)29 are the most active for hydrogen evolution/oxidation19,29 and hydrogen/deuterium exchange at acidic conditions.30 Brønsted-Evans-Polanyi (BEP) relations suggest that the activation energy for any elementary step varies linearly with its energy of reaction, hence, a weaker bonding of H adatoms facilitates their addition to the reacting substrate.31,32 The binding strength, however, also influences the coverage of hydrogen through the adsorption equilibrium. We probe this hypothesis here, showing the hydrogen binding energy to be adjusted by the pH, leading to an increase of the rate of phenol hydrogenation. The rate of hydrogenation of phenol on Pt at high H2 pressure increases 15-fold with decreasing pH. Based on experimentally determined changes in HBE with pH6,7,9–11,16–18 this increase in activity is attributed to a decrease in the adsorption energy of the reactant (hydrogen adatoms) and the decrease in activation energy for its addition to adsorbed phenol, as expected for a BEP relation with a slope of 1. This supports the model proposed that the HBE change with pH is responsible for the electrochemical HER rate dependence on pH (with a BEP slope of 0.8 for Pt/C).6 We show that under the conditions investigated, Rh does not have the same dependence on pH as Pt for hydrogenation, and attribute this to the qualitative difference

ACS Paragon Plus Environment

ACS Catalysis 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 2 of 9

in hydrogen coverage on Pt versus Rh under similar reaction conditions. We also present ab initio molecular dynamics (AIMD) studies on Pt(110) showing that the increase in hydrogen adsorption strength with increasing pH is related to the effect of both water and ions (including OH- ions) on the interfacial electrostatic potential gradient (i.e., dipole) in the double layer. Although it is currently unknown, which Pt facet has the highest activity for phenol hydrogenation, the Pt(110) surface facet has been selected as the model surface since the HER activity has been previously shown to follow the order (110) > (100) > (111).12,26,33–35 Several modeling studies have attempted to build methodology capable of incorporating pH variation,20,36–38 but the studies use finite, static water layers to simulate the interface,36 or small unit cell sizes36. In a recent computational study by Hansen and co-workers, explicit water/Pt(111) simulations show that the relative coverage of adsorbed species, including Had and OHad, significantly vary from the relative coverage on a bare Pt(111) surface or a surface modeled with a static water layer.38 In this work, we conducted AIMD simulations to provide insight into the impact of electrolyte ion distribution and interfacial solvent reorganization on the HBE. The simulations used a large simulation cell (~800 atoms) to avoid substantial impact of artificial boundary conditions,39 and explicitly included multiple ion pairs to simulate a variable pH environment. The simulations were performed at temperatures consistent with the upper end of the experimental temperature range, to allow for statistically relevant sampling of water, known to require long relaxation times (~5 ps at 298 K).40

KHCO3, pH 9.2). Non-zero conversions at 0 minutes for pH 2.5 and 5 are due to reaction during heat-up of the reactor under these conditions. The heat-up times and % conversion at 0 minutes were reproducible for a given pH value.

RESULTS

Figure 2. Initial TOFs of phenol hydrogenation for 5 wt% Pt/C at 80°C and 20 bar H2, and HBEs for Pt(110) plotted as a function of pH. TOFs are calculated from the initial slopes in Figure 1. HBEs are relative to 1 bar H2,g at given pH (kJ per mole H adatom) from cyclic voltammetry at room temperature by Sheng et al.17 The difference in experimental HBE at pH 1 and 3 is due to variation in the experimental measurement in different electrolytes from reference 17.

Phenol hydrogenation on Pt. Figure 1 shows that the rate of phenol conversion on Pt/C is pH-dependent at 20 bar H2 and 80 °C (where the order in H2 is zero at pH 54). The reaction rate at low pH was much higher than at moderate or high pH. No conversion was seen at pH 9.2, and extrapolating the rate to this pH puts conversion below the detection limit.

The initial turnover frequency (TOF, product molecule per metal surface atom per time) at 80°C increases with decreasing pH as shown in Figure 2, which also shows the HBE for H adatoms on Pt(110) measured by Sheng et al.17 from electrochemical cyclic voltammetry as a function of pH. The HBEs show a trend towards weaker hydrogen binding energy (less negative value) at lower pH.

Table 1. Hydrogen binding energy (HBE) (in kJ per mol of Had) and change in the interfacial electrostatic potential energy (Δφ) (in kJ mol-1) calculated using the lowest energy structures for hydrogen adsorption on the Pt (110) surface in acidic, water and alkaline solution.

Figure 1. Catalytic hydrogenation of phenol (172 mM in water) on Pt/C at 20 bar H2 and 80 °C at different pH in an batch reactor. Conversion of ~40% was observed at 90 minutes for pH 8, shown by the gray dotted line. Conversion at 232 minutes was zero for the carbonate/bicarbonate buffer solution (10 mM K2CO3/90 mM

pH

HBE (kJ mol-1)

Δφ (kJ mol-1)

0

−39.1

−83.9

7

−41.6

−79.1

14

−50.7

−65.6

Table 1 reports our theoretically calculated HBE on Pt (110) relative to gas-phase H2 as a function of pH. The Pt (110) surfaces are maintained neutral in the simulations to emulate thermal catalytic reaction conditions and observe changes in the HBE primarily with the solution pH. The trend also indicates that hydrogen binding increases with pH with an energy shift of approximately 12 kJ mol-1 across the pH scale. A corresponding 18 kJ mol-1 shift is observed in the electrostatic potential energy drop across the interfacial region (Δφ). We rationalize that the differences in the solvent structure and composition impact the local electric fields

ACS Paragon Plus Environment

Page 3 of 9 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis in/near the interfacial region, which leads to shifts in Δφ, that further manifests as a change in HBE with pH (see below). Compared to the pH-dependent values of HBE extrapolated from cyclic voltammograms reported by Sheng15 and Koper,14 the simulations performed under TCH conditions give the same trends (~0.9 kJ mol-1 stronger HBE with increase of pH by one). As the change and not the absolute magnitude of the HBE with pH is central in the discussion, we base the analysis on the experimentally determined HBE from Sheng et al,17 but the results are similar, if we use the calculated values in Table 1 with an offset of approximately 25 kJ mol-1. Figure 3 shows that ln(TOF) for phenol hydrogenation scales with the reported HBE.17 The slope is approximately 1 (and would also be ~1 if we used the HBE values from Table 1, although the intercept would change). If we extrapolate the trend line to the HBE/RT at pH 9.2, it leads to a TOF of approximately 1.5 s-1, which would only produce immeasurably low cyclohexanone/cyclohexanol concentrations, in line with the measured inactivity at pH 9.2 in Figures 1 and 2. We note here that the log(TOF) also scales with log[H+] (Figure S1). At 1 bar H2 and room temperature, the rate of catalytic hydrogenation of phenol on Pt did not increase with decreasing pH (Figure S2, Table S1), unlike at 20 bar H2 and 80 °C. The selectivity to cyclohexanone versus cyclohexanol is also unaffected by a change of pH from 1 to 5 here (Figure S2). Apart from the difference in the effect of pH on rates at the two conditions in Figure 2, the rate depends on the H2 pressure near 1 bar (first order for 5-60 °C),4 unlike what is observed at 20 bar, where the hydrogenation is zero-order in hydrogen at 60-100 °C.4 In both cases, the order in phenol is also zero.4 Based on these orders, the hydrogen coverage is low at 1 bar and 25 °C on Pt/C, and higher at 20 bar H2 and 80°C. Zero order in both phenol and hydrogen as observed at 20 bar H2 and 80°C generally occurs when both reacting species are at similar high coverages (assuming a competitive Langmuir-Hinshelwood surface reaction). The difference in the effect of pH in Figure 2 and Figure S2 is attributed to this quantitative difference in hydrogen coverage within the same BEP relation that explains Figure 3.

Figure 3. pH-dependent phenol hydrogenation rate at 80°C and 20 bar H2 and HBE data of Figure 2 replotted as ln(TOF) for

hydrogenation of phenol on 5 wt% Pt/C versus HBE/RT, using the HBE reported on Pt(110),17 where R is the ideal gas constant and T is 80 °C (353 K). The best-fit straight line is also shown.

Phenol hydrogenation on Rh. Rh shows the same decrease in HBE as a function of pH as Pt, although with 1-2 kJ mol-1 weaker HBE than Pt,6 and so may be expected to behave similarly to Pt. Like Pt/C, Rh/C at 1 bar H2 did not show a change in phenol hydrogenation rates with pH from 1 to 5. However, at the conditions of Figures 1-3 (20 bar H2 and 80 °C), Rh/C was similarly active at pH 1 and 8 (Table 2), unlike Pt/C. Even at 70 bar H2 and 40 °C (in an attempt to increase the coverage of H) this same independence of reaction rates on pH was observed (Table 2). Rh/C showed a first-order increase in reaction rate with H2 pressure up to 70 bar at 40 °C and pH 5 (Figure S5b,) as well as at 80 °C and pH 1 (Figure S5a). This implies Rh/C has a lower hydrogen coverage than Pt/C at >20 bar. We hypothesize that the first-order rate dependence is less caused by the slightly lower HBE on Rh17,41, but rather by the higher phenol (and also benzene) binding energy on Rh than on Pt.42– 45

Table 2. Hydrogenation rates for 5 wt% Rh/C catalyst. Starting concentration of phenol is 18 mM in aqueous solution at room temperature, 172 mM phenol at 80 °C (phenol is zero-order for this reaction). The pH is adjusted using HClO4, bicarbonate/carbonate or KOH. HBE are reported from the hydrogen underpotential desorption (Hupd) peak potential from reference 6, and so are an uncorrected estimate for HBE in different electrolytes. Conversion vs. time data are in Figures S4-S5. pH

HBE on Rh represented by Hupd CV peak (V vs. RHE)6

Temp (°C)

Pressure (bar)

TOF (s-1)

1

0.1

25

1

0.11

2.5

0.125

25

1

0.09

5

0.13

25

1

0.12

1

0.1

80

20

25.2

8

0.16

80

20

20.5

1

0.1

40

70

3.09

8

0.16

40

70

3.83

5

0.13

40

20

0.78

5

0.13

40

30

1.64

5

0.13

40

50

2.02

Calculated impact of pH on Pt(110) surface in solution. Figure 4 shows the density profiles for hydrogen, water and its ions, and Na+/Cl- ions in acidic (HCl), neutral and alkaline (NaOH) media computed from AIMD simulations near the Pt(110) surface. In contrast to the similar position and width of the peak in H, the interfacial distribution of the water molecules is distinctly different in acidic, neutral and alkaline solutions (Figure 4). While in equilibrium the thermodynamic activity of water is identical across the reactor, the

ACS Paragon Plus Environment

ACS Catalysis 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

concentration of water molecules in the first solvent layer (0 < z < 2.6 Å) is lowered in non-neutral solutions. This decrease is attributed to the higher concentration of electrolyte ions in the interfacial region (0 < z < 4 Å) as shown in Figure 4. The presence of ions in the double layer, either Cl− ions in acidic conditions (Figure 4a) or Na+ and adsorbed OH− ions in alkaline conditions (Figure 4c), disrupts the distribution and packing of the water molecules near the Pt (110) surface (see Figures 4 and S6). The presence of ions in the interfacial region also modifies the extent of polarization in the water molecules. Figure S7 shows the orientation of the water dipoles in the first solvent layer. The average dipole angle of the first solvent layer with respect to the surface normal was calculated to be 34ο in acidic, 36ο in neutral but 21ο in alkaline solution. This implies that the presence of adsorbed OH− and Na+ perturbs the orientation of the water, hence the hydrogen of H2O is tilted more towards the surface. We rationalize this as a combined effect of charge transfer from OH- to the surface (see Figure S8) and water solvation of Na+ near the surface. Figure S8 provides evidence of the impact of ions at the interface on the charged state of the metal. Under alkaline conditions, the presence of Na+ and OH− ions at the interface makes the Pt(110) surface more negative relative to the Pt(110) surface under acidic conditions. We attribute this to the charge transfer due to adsorbed OH- at the surface. This is reflected in the electrostatic drop showing lower Δφalkaline (~ −65.6 kJ mol-1) than Δφacidic (~ −83.9 kJ mol-1) and a corresponding increase in HBE (see Table 1). In a neutral environment, a drop of Δφneutral ~ −79.1 kJ mol-1 indicates that a well-ordered solvent alone is unlikely to create significant shifts in the metal work function. These observations signify that the presence of electrolytic ions in the interfacial region play a crucial role in determining the adsorption/desorption chemistry of H2.

Figure 4. Atomic density profiles (in molecules/Å3) showing distribution of species as a function of the Pt (110) surface normal direction (z) in (a) acidic solution (pH = 0), (b) water (pH = 7) and, (c) basic solution (pH = 14). z = 0 denotes the metal surface. Values were obtained using atom number distributions averaged over all frames of the AIMD NVT trajectory. Density profile color representations: Pt (gray), surface adsorbed H (red), water oxygens OH2O (blue), hydronium oxygens OH3O+ (brown), Cl− ions (green), hydroxide oxygens OOH− (orange) and Na+ ions (purple).

DISCUSSION Phenol hydrogenation on Pt at high hydrogen coverage. The rate of hydrogenation on Pt was affected by the pH (Figures 1, 2), with rates typically increasing with increasing H3O+ concentrations. At 20 bar and at pH 5 from 60-100 °C, the reaction was zero-order in both phenol and H2 for Pt/C.4 Thus, changing the equilibrium constant of adsorption for either species by small factors would not affect the coverage of phenol and hydrogen. The change in rate with hydronium

ACS Paragon Plus Environment

Page 4 of 9

Page 5 of 9 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis ion concentration is hypothesized to be caused by the change in the hydrogen binding energy. We assume that the rate-determining step for phenol hydrogenation is the first addition of an adsorbed hydrogen atom to an adsorbed phenol, consistent with a broad range of kinetic measurements on Pt/C (and fewer measurements on Rh/C).4,46 This proposed mechanism is consistent with the kinetic model for this reaction in our earlier work,4 whereby surface hydrogen addition is rate-limiting. It was proposed there partially because the dependences of the rate on partial pressures and temperature were very similar to those explained by the same mechanism in a well-established model for the kinetics of gas-phase hydrogenation of aromatics over Pd developed by Vannice.47 The rate of reaction is thus expected to be: ― 𝐸𝑎,𝑒𝑓𝑓 ― 𝐸𝑎 1 𝑇𝑂𝐹 = 𝐴exp = 𝐵𝜃 𝜃 exp

(

𝑅𝑇

)

𝐻 𝑃ℎ

( ) 𝑅𝑇

where Ea,eff is the effective activation barrier including temperature dependence of coverages, Ea is the activation barrier, A is the preexponential factor, B is an independent prefactor including the activation entropy, 𝜃𝐻 and 𝜃𝑃ℎ are the coverages of adsorbed hydrogen and adsorbed phenol, respectively. When the rate is zero-order in both phenol and hydrogen as in Figures 1-3, the product of the coverages must be constant, so 𝐴 ≡ 𝐵𝜃𝐻𝜃𝑃ℎ is constant. We will neglect here any effect of coverages on Ea. Scheme 1 shows the enthalpy diagram of the rate-determining step under three values of pH. Phenol* represents either adsorbed phenol or a partially hydrogenated intermediate, whichever one is involved in the rate-determining H addition. A less negative H binding energy with decreasing pH raises the total energy of adsorbed reactants. Scheme 1. Enthalpy diagram of a hydrogenation step at different pH values. An endothermic reaction step is depicted but not necessary for this analysis.

The reaction energy of the rate-determining hydrogenation step shown in Scheme 1 is: ∆𝐻𝑟𝑑𝑠 = ∆𝐻0𝑓,𝑃ℎ ― 𝐻 ― ∆𝐻0𝑓,𝑃ℎ ― ∆𝐻0𝑓,𝐻 ∗

2

= ∆𝐻0𝑓,𝑃ℎ ― 𝐻 ― ∆𝐻0𝑓,𝑃ℎ ― 𝐻𝐵𝐸

In addition to the effect of pH on the HBE, the pH may also affect the adsorption enthalpy of phenol, but it is likely that both the adsorbed phenol reactant and its hydrogenated product in this step will be affected to a similar extent, so any energy change with pH is expected to cancel following the above analysis. Therefore, ∆𝐻0𝑓,𝑃ℎ ― 𝐻 ― ∆𝐻0𝑓,𝑃ℎ is assumed to be constant with pH. Thus, at higher pH (stronger hydrogen binding, or more negative HBE), the reaction energy of the rate-determining step would be less thermodynamically favorable. According to BEP relations, the activation energy

of the rate-determining step varies linearly with its reaction enthalpy: 𝐸𝑎,𝑒𝑓𝑓 = 𝛼 + 𝛽∆𝐻𝑟𝑑𝑠 3 where β is constant between 0 and 1, α is a constant, and Ea is ~34 kJ mol-1 at 20 bar and pH 5 on Pt.4 Combining this with Equation 1 leads to an expression for the turnover frequency (see Supporting Information for derivation):

(

𝑇𝑂𝐹 = 𝐴′𝑒𝑥𝑝

)

𝛽 × 𝐻𝐵𝐸 𝑅𝑇

4

where A’ is a new constant, differing from A. This can be rewritten as:

( )

ln (𝑇𝑂𝐹) = ln (𝐴′) + 𝛽

𝐻𝐵𝐸 𝑅𝑇

5

Thus, ln(TOF) should increase linearly with HBE (weaker binding of Had to the surface). Based on the slope of ln(TOF) vs. HBE/RT for Pt(110) from Figure 3, β is 1.0, the maximum for BEP relations, which points to a late transition state for the rate-determining step. Thus, the transition state resembles the product with respect to the HBE, i.e., the H-Pt bond is nearly completely broken and a strong C-H bond is already present. In contrast, the dehydrogenation of sp2 and sp3 hybridized C-H bonds in adsorbed cyclic hydrocarbons on Pt(111) proceeds with transition states that resembles the products, i.e., where the C-H bond is nearly fully broken, and strong H-Pt bonding is already present.31 For the reverse reaction, i.e., hydrogenation, the transition state should then resemble the reactants. This opposite behavior in vacuum and aqueous phase is tentatively attributed to the impact of the aqueous environment on the transition state, or to the fact that these vacuum reactions in reverse would be adding H to an unstable intermediate hydrocarbon that has an odd number of H atoms, whereas here we are adding H to a stable adsorbed molecule. This analysis strongly suggests that the change in HBE with pH is at least partly responsible for the increase in rate of hydrogenation on Pt at high hydrogen coverages. Furthermore, AIMD simulations show that the adsorption is significantly impacted by the presence of ions (including OH- ions) and the corresponding orientation of water at the interface. For a more detailed discussion of the recent literature, see the Supporting Information. Phenol hydrogenation on Pt and Rh described by BEP relations at low hydrogen coverage. For Pt/C, a positive reaction order in H2 was observed at pressures below 10 bar,4 indicating that hydrogen coverages are significantly below 1. Weakening the HBE by lowering the pH has been argued to decrease the hydrogen coverage further, thus, cancelling the effect of pH on the barrier (Equation 3) in the net effect on rate. Under low H* coverage conditions, Equation 1 can be written as (see Supporting Information): (𝛽 ― 1)𝐻𝐵𝐸 6 𝑇𝑂𝐹 = 𝐶′𝑃𝑛𝐻2𝑒𝑥𝑝 𝑅𝑇

(

)

where n = ½ or 1 (depending on whether the dominant removal pathway for adsorbed H is by H2 desorption or reaction with an organic).

ACS Paragon Plus Environment

ACS Catalysis 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

From Equation 6, it is evident that under conditions where the hydrogen coverage is low (H2 order is 1/2 to 1 based on this simple model), the TOF would be independent of the HBE if β is 1. This agrees well with the fact that the hydrogenation rates on Pt at 1 bar H2 is independent of pH, whereas it strongly depends on pH at 20 bar H2, when the reaction order in H2 is zero. Assuming that phenol hydrogenation on Rh/C has the same mechanism as proposed above for Pt/C, Equation 6 also describes the independence of Rh-catalyzed hydrogenation on pH, having also has a positive order in H2 under the conditions tested here (Figure S5, Table 2), consistent with low H* coverage. Based on the measured HBEs for other metals, it is expected that Pd, which has a HBE greater than Pt,41 would behave similarly to Pt. Ir, which has a HBE even lower than Rh41, would behave as Rh. Role of hydroxide at Pt solid/liquid interface in alkaline conditions. From our computational results, under alkaline conditions, adsorbed OH species are at the Pt(110) surface at low concentrations (~0.05 ML) when the pH is high, and are partially responsible for the observed stronger HBE in alkaline conditions compared to acidic conditions. These results are consistent with observations that Pt-H and Pt-OH may co-exist under alkaline conditions, summarized in the next paragraph. Using in operando ambient pressure X-ray photoelectron spectroscopy in alkaline conditions, Stoerzinger et al.48 identify Pt-oxides and adsorbed OH on the surface as a function of applied potential and report coexistence of Pt-H and Pt-OH species. Note, however, that the observed surface coverage of OHad does not change with potential in a way consistent with a recently proposed model 14,18,26,49 whereby OH adsorption, simultaneous with removal of 2 Had to make 2 H+aq, is responsible for CV peaks seen in the Hupd region that have traditionally been attributed to simple desorption of Had to make H+aq. Instead, it changes the opposite direction with potential. By comparing experimental and DFT-derived XANES spectra, Jia and co-workers50 corroborated the presence of adsorbed OHad, albeit on Ru and PtRu surfaces. Since Ru is much more oxophillic than Pt, this does not prove that such species are present on pure Pt surfaces. Our results here indicate that although adsorbed OH may be present at the surface at low concentrations in the Hupd CV regime, the change in HBE with pH is what impacts the rate of thermal hydrogenation. Role of pH on hydrogen binding energy. Since the effect of pH on the HBE plays such an important role in the hydrogenation rate of phenol in aqueous solution at high H* coverages, and in the HER reaction, it is important to mention other explanations for this effect besides the model that comes out of the current calculations. A recent perspective discusses the different hypotheses in detail, and hypothesizes that the change in structure of interfacial water with pH affects the apparent hydrogen binding energy measured by cyclic voltammetry.16 Recent Quantum Mechanics Molecular Dynamics calculations by Goddard et al. 20 on the Pt(100) surface attribute the hydrogen binding energy change to movement of interfacial water further from the Pt as the potential shifts negative.16 Although the measurement of the HBE is referenced to RHE, the potential vs. SHE (at any fixed potential vs. RHE) is different for different pH values, due to changes in the electric field of the electrochemical double layer.6,51 With increasing pH, the equilibrium potential (i.e., 0 V vs. RHE) shifts vs. SHE

by 59 mV per pH unit (from the Nernst equation). This shift can only occur if a dipole layer is added (in the double layer) that changes the electric field to match the voltage shift, oriented such that the positive ends of the dipole are towards the solution (away from metal surface). This change in the double layer with pH could affect the energy of adsorbed species on the surface, and the orientation of ions and water molecules near the surface. The orientation of interfacial water on a gold electrode is known to change with applied potential (no change in pH, but change with respect to SHE).52 Similar changes are also observed on Ag(100),53 and so it may be expected that the change in pH could also result in changes in the orientation of water, similar to what is computed in Figure S7 for Pt(110). In this simple model, an adsorbate that (in vacuum) causes the metal work function to decrease would form a dipole (consisting of the metal and the adsorbed species) with the positive end away from the metal, which would be stabilized by the change in the double layer field that arises from increasing pH. Hydrogen adatoms on metal surfaces such as platinum form a dipole with the positive end at the hydrogen,54,55 and so changes in the double layer could influence its binding energy via dipole-dipole interactions (as attractions at higher pH, making it bind more strongly).

CONCLUSIONS The impact of the hydronium ion concentration on the catalytic hydrogenation on Pt is attributed to the change in the hydrogen binding energy. At high coverage of hydrogen adatoms, the weakening of the hydrogen binding energy at lower pH decreases the activation energy for the ratedetermining step by the same amount as the HBE, thus, increasing the rate. At low coverages of hydrogen, the weakening of the hydrogen binding still decreases the intrinsic activation energy of the rate-determining step but additionally decreases the hydrogen coverage as well. This compensates the positive effect resulting in a rate independent of the hydronium ion concentration. The significantly weaker HBE in water requires high H2 pressures to achieve higher H coverages. Our calculations corroborate the weakening of the HBE at lower pH values.

METHODS Catalysis. Phenol (≥99% purity) was added to 80 mL of Millipore water to reach the desired phenol concentration. Perchloric acid, potassium carbonate/bicarbonate or potassium hydroxide was added to vary the pH (verified by pH strips with accuracy of +/- 0.25 pH units). Platinum or rhodium (Pt/C or Rh/C, 5 wt% metal) was added to the reactor (approximately 10 mg). Characterization of Pt/C and Rh/C is discussed in previous work.4 H2 was introduced to the system below the liquid surface at 30 bar, then evacuated, repeated three times to ensure removal of air from the reactor, then introduced to reach the desired H2 reaction pressure. During heating, the stirring rate was kept at a very low value, 40 rpm. Once 80 °C was reached, the stirring rate was increased to 700 rpm. For Rh/C, 40 °C was used to measure rates to allow for longer times without complete conversion at high pressure. Turnover frequencies (TOFs) were calculated based on the rate of conversion once the reaction temperature was reached, normalized to number of sites. The number of sites of the catalyst were determined from H2 gas chemisorption. For 1 bar rate measurements, phenol was added to 50 mL of Millipore

ACS Paragon Plus Environment

Page 6 of 9

Page 7 of 9 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis water and perchloric acid was used to vary the pH. H2 was bubbled through the glassware reactor while it was stirred using a stir bar, and 1 mL aliquots were taken periodically to measure the conversion as a function of time. Chemicals were purchased from Sigma-Aldrich. We have conducted several studies to verify that intrinsic kinetics are being measured (as outlined in reference 4), including stir rate dependence to ensure no mass transfer limitations, and measurement of rates as a function of temperature, with much higher rates at higher temperature as shown in reference 4. The higher rates at higher temperature and corresponding barriers imply no diffusion limitation. Unfortunately, due to the particle size dependence on rate reported,56 testing the TOF by varying the catalyst loading, which affects the particle size, was not useful to ensure that we are operating under kinetic limitations. The absence of internal diffusion limitation has been indicated for this catalyst in previous work in reference 3 by calculating the Thiele modulus and Weisz modulus. Product analysis. Product quantification was carried out by gas chromatograph with attached mass spectrometer. The products were extracted using dichloromethane and an internal dimethoxybenzene standard was added to compare to standards prepared with known concentrations for quantification. Computational methods and details. Periodic densityfunctional-theory (DFT) based ab initio molecular dynamics (AIMD) simulations were performed with the Perdew, Burke and Ernzerhoff (PBE)57 functional as implemented in the CP2K package.58,59 Valence wave functions were expanded with double-zeta quality Gaussian basis sets.60 Core electrons were represented using norm-conserving pseudopotentials.61 Electrostatic terms were calculated using an additional auxiliary plane-wave basis set with a 400 Ry cutoff. Grimme’s DFT-D362 corrections were applied to account for dispersion forces or van der Waals interactions. The Brillouin zone was integrated using a Γ-point approximation. A Pt (110) surface slab of cell dimensions 23.493 × 16.612 × 26.668 Å was cleaved and consisted of 180 Pt atoms (45 atoms/layer) with 20 Å vacuum space. The low index (110) facet was chosen due to previous evidence of non-Nernstian, pH induced shifts in cyclic voltammograms peaks in the hydrogen adsorption region across the Pt(110) and Pt(110) steps.26,63 To simulate pH variation, three distinct models that varied in the interfacial solvent composition were created: (a) 25 H2 + 180 H2O + 6 HCl, (b) 25 H2 + 180 H2O and (c) 25 H2 + 180 H2O + 6 NaOH were placed above the Pt (110) surface, respectively. The number of water molecules was chosen to reflect a bulk water density of 1.3 gm/cm3 with a H2 partial pressure (pH2) of 6700 bar. Figure S9 provides a structural representation of the models. AIMD simulations were performed in the canonical ensemble (NVT) at 373 K to obtain statistics on the dynamic distribution of the solvent molecules. A Nosé-Hoover thermostat with a 1 fs timestep was used. For analysis, 4-6 ps of data were collected from the AIMD trajectories post equilibration. A time-averaged concentration of H3O+ and OH− species was calculated (see Table S2), corresponding to pH values of 0, 7 and 14 in the three models, respectively. In addition, the NVT simulations were followed by simulated temperature annealing to 0 K to obtain optimized structures to obtain the hydrogen binding energy (HBE) and the interfacial electrostatic

potential energy drop (Δφ), as reported in Table 1. The HBE was calculated using Equation 7, where H* denotes surface adsorbed hydrogen, Sol denotes solvent consisting of H2, H2O/H3O+/OH-, and Na+/Cl-, and Tot = Pt (110) + Sol + H*:

𝐻𝐵𝐸 =

𝐷𝐹𝑇 𝐷𝐹𝑇 𝐷𝐹𝑇 𝐷𝐹𝑇 𝐷𝐹𝑇 𝐸𝐷𝐹𝑇 𝑇𝑜𝑡 ― 𝐸𝑇𝑜𝑡 ― 𝐻 ∗ ― 𝑛𝐻 ∗ ∙ 𝐸𝐻2 /2 ― (𝐸𝑆𝑜𝑙 ― 𝐸𝑆𝑜𝑙 ― 𝐻 ∗ ― 𝐸𝐻 ∗ )

𝑛𝐻 ∗

7

ASSOCIATED CONTENT Supporting Information. Additional reaction rate data, rate equation model derivation, discussion of hydrogen coverage and computational details. This material is available free of charge via the Internet at http://pubs.acs.org.

AUTHOR INFORMATION Corresponding Authors * Email: [email protected], [email protected], [email protected]

Author Contributions NS completed the measurements at >1 bar pressure. NS and GC completed 1 bar pressure measurements. MSL and SA completed the computational work. All authors have given approval to the final version of the manuscript. CTC helped interpret the data and write the paper.

ABBREVIATIONS HBE, hydrogen binding energy; HER, hydrogen evolution reaction; HOR, hydrogen oxidation reaction; BEP, BrønstedEvans-Polanyi; TOF, turnover frequency.

ACKNOWLEDGMENTS The research described in this paper is part of the Chemical Transformation Initiative at Pacific Northwest National Laboratory (PNNL), conducted under the Laboratory Directed Research and Development Program at PNNL, a multiprogram national laboratory operated by Battelle for the U.S. Department of Energy. N.S. acknowledges the Washington Research Foundation Innovation Fellowship. C.T.C. acknowledges the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences, Chemical Sciences Division Grant No. DE-FG02-96ER14630 for support of this work.

REFERENCES (1)

(2) (3)

(4)

(5)

Singh, N.; Nguyen, M.-T.; Cantu, D. C.; Mehdi, B. L.; Browning, N. D.; Fulton, J. L.; Zheng, J.; Balasubramanian, M.; Gutiérrez, O. Y.; Glezakou, V.-A.; Rousseau, R.; Govind, N.; Camaioni, D. M.; Campbell, C. T.; Lercher, J. A. Carbon-Supported Pt during Aqueous Phenol Hydrogenation with and without Applied Electrical Potential: X-Ray Absorption and Theoretical Studies of Structure and Adsorbates. J. Catal. 2018, 368, 8–19. Song, Y.; Chia, S. H.; Sanyal, U.; Gutiérrez, O. Y.; Lercher, J. A. Integrated Electrocatalytic Conversion of Substituted Phenols and Diaryl Ethers. J. Catal. 2016, 344, 263–272. Song, Y.; Sanyal, U.; Pangotra, D.; Camaioni, D. M.; Gutiérrez, O. Y.; Johannes A. Lercher; Holladay, J.; Camaioni, D. M.; Gutiérrez, O. Y.; Lercher, J. A. Hydrogenation of Benzaldehyde via Electrocatalysis and Thermal Catalysis on Carbon-Supported Metals. J. Catal. 2018, 359, 68–75. Singh, N.; Song, Y.; Gutiérrez, O. Y.; Camaioni, D. M.; Campbell, C. T.; Lercher, J. A. Electrocatalytic Hydrogenation of Phenol over Platinum and Rhodium: Unexpected Temperature Effects Resolved. ACS Catal. 2016, 6, 7466–7470. Sheng, W.; Gasteiger, H. A.; Shao-Horn, Y. Hydrogen Oxidation and Evolution Reaction Kinetics on Platinum: Acid vs Alkaline Electrolytes. J. Electrochem. Soc. 2010, 157, B1529.

ACS Paragon Plus Environment

ACS Catalysis (6)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

(7)

(8)

(9)

(10)

(11)

(12) (13)

(14)

(15)

(16)

(17)

(18)

(19) (20)

(21)

(22) (23)

(24)

Zheng, J.; Sheng, W.; Zhuang, Z.; Xu, B.; Yan, Y. Universal Dependence of Hydrogen Oxidation and Evolution Reaction Activity of Platinum-Group Metals on pH and Hydrogen Binding Energy. Sci. Adv. 2016, 2, 1–8. Durst, J.; Siebel, A.; Simon, C.; Hasché, F.; Herranz, J.; Gasteiger, H. A. New Insights into the Electrochemical Hydrogen Oxidation and Evolution Reaction Mechanism. Energy Environ. Sci. 2014, 7, 2255. Zheng, J.; Zhou, S.; Gu, S.; Xu, B.; Yan, Y. Size-Dependent Hydrogen Oxidation and Evolution Activities on Supported Palladium Nanoparticles in Acid and Base. J. Electrochem. Soc. 2016, 163, F499–F506. Bhowmik, T.; Kundu, M. K.; Barman, S. Palladium NanoparticleGraphitic Carbon Nitride Porous Synergistic Catalyst for Hydrogen Evolution/Oxidation Reactions over a Broad Range of pH and Correlation of Its Catalytic Activity with Measured Hydrogen Binding Energy. ACS Catal. 2016, 6, 1929–1941. Marković, N. M.; Grgur, B. N.; Ross, P. N. TemperatureDependent Hydrogen Electrochemistry on Platinum Low-Index Single-Crystal Surfaces in Acid Solutions. J. Phys. Chem. B 1997, 101, 5405–5413. Marković, N. M.; Sarraf, S. T.; Gasteiger, H. A.; Ross, P. N. (Jr. . Hydrogen Electrochemistry on Platinum Low-Index SingleCrystal Surfaces in Alkaline Solution. J. Chem. Soc. Faraday Trans. 1996, 92, 3719–3725. Strmcnik, D.; Lopes, P. P.; Genorio, B.; Stamenkovic, V. R.; Markovic, N. M. Design Principles for Hydrogen Evolution Reaction Catalyst Materials. Nano Energy 2016, 29, 29–36. Strmcnik, D.; Kodama, K.; Van Der Vliet, D.; Greeley, J.; Stamenkovic, V. R.; Marković, N. M. The Role of Non-Covalent Interactions in Electrocatalytic Fuel-Cell Reactions on Platinum. Nat. Chem. 2009, 1, 466–472. Chen, X.; McCrum, I. T.; Schwarz, K. A.; Janik, M. J.; Koper, M. T. M. Co-Adsorption of Cations as the Cause of the Apparent pH Dependence of Hydrogen Adsorption on a Stepped Platinum Single-Crystal Electrode. Angew. Chemie - Int. Ed. 2017, 56, 15025–15029. Sheng, W.; Myint, M.; Chen, J. G.; Yan, Y. Correlating the Hydrogen Evolution Reaction Activity in Alkaline Electrolytes with the Hydrogen Binding Energy on Monometallic Surfaces. Energy Environ. Sci. 2013, 6, 1509–1512. Zheng, J.; Nash, J.; Xu, B.; Yan, Y. Towards Establishing Apparent Hydrogen Binding Energy as the Descriptor for Hydrogen Oxidation/Evolution Reactions. J. Electrochem. Soc. 2018, 165, H27–H29. Sheng, W.; Zhuang, Z.; Gao, M.; Zheng, J.; Chen, J. G.; Yan, Y. Correlating Hydrogen Oxidation and Evolution Activity on Platinum at Different pH with Measured Hydrogen Binding Energy. Nat. Commun. 2015, 6, 5848. van Der Niet, M. J. T. C.; Garcia-Araez, N.; Hernández, J.; Feliu, J. M.; Koper, M. T. M. Water Dissociation on Well-Defined Platinum Surfaces: The Electrochemical Perspective. Catal. Today 2013, 202, 105–113. Parsons, R. The Rate of Electrolytic Hydrogen Evolution and the Heat of Adsorption of Hydrogen. Trans. Faraday Soc. 1957, 54, 1053–1063. Cheng, T.; Wang, L.; Merinov, B. V.; Goddard, W. A. Explanation of Dramatic pH-Dependence of Hydrogen Binding on Noble Metal Electrode: Greatly Weakened Water Adsorption at High pH. J. Am. Chem. Soc. 2018, 140, 7787–7790. Zheng, Y.; Jiao, Y.; Qiao, S.; Vasileff, A. Hydrogen Evolution Reaction in Alkaline Solution: From Theory, Single Crystal Models, to Practical Electrocatalysts. Angew. Chemie Int. Ed. 2018, 57, 7568–7579. Zuo, X. Q.; Chen, W.; Yu, A.; Le Xu, M.; Cai, J.; Chen, Y. X. pH Effect on Acetate Adsorption at Pt(111) Electrode. Electrochem. commun. 2018, 89, 6–9. McCrum, I. T.; Chen, X.; Schwarz, K. A.; Janik, M. J.; Koper, M. T. M. Effect of Step Density and Orientation on the Apparent pH Dependence of Hydrogen and Hydroxide Adsorption on Stepped Platinum Surfaces. J. Phys. Chem. C 2018, 122, 16756–16764. He, Z. Da; Hanselman, S.; Chen, Y. X.; Koper, M. T. M.; CalleVallejo, F. Importance of Solvation for the Accurate Prediction of Oxygen Reduction Activities of Pt-Based Electrocatalysts. J. Phys. Chem. Lett. 2017, 8, 2243–2246.

(25) (26) (27)

(28)

(29) (30)

(31)

(32) (33) (34)

(35)

(36) (37)

(38) (39)

(40)

(41) (42) (43)

(44)

(45)

Page 8 of 9 Wang, H.; Gao, L. Recent Developments in Electrochemical Hydrogen Evolution Reaction. Curr. Opin. Electrochem. 2018, 7, 7–14. McCrum, I. T.; Janik, M. J. pH and Alkali Cation Effects on the Pt Cyclic Voltammogram Explained Using Density Functional Theory. J. Phys. Chem. C 2016, 120, 457–471. Xue, S.; Garlyyev, B.; Watzele, S.; Liang, Y.; Fichtner, J.; Pohl, M. D.; Bandarenka, A. S. Influence of Alkali Metal Cations on the Hydrogen Evolution Reaction Activity of Pt, Ir, Au, and Ag Electrodes in Alkaline Electrolytes. ChemElectroChem 2018, 4, 1–5. Garlyyev, B.; Xue, S.; Watzele, S.; Scieszka, D.; Bandarenka, A. S. Influence of the Nature of the Alkali Metal Cations on the Electrical Double-Layer Capacitance of Model Pt(111) and Au(111) Electrodes. J. Phys. Chem. Lett. 2018, 9, 1927–1930. Nørskov, J. K.; Bligaard, T.; Logadottir, A.; Kitchin, J. R.; Chen, J. G.; Pandelov, S.; Stimming, U. Trends in the Exchange Current for Hydrogen Evolution. J. Electrochem. Soc. 2005, 152, J23. Singh, N.; Upham, D. C.; Metiu, H.; McFarland, E. W. Gas-Phase Chemistry to Understand Electrochemical Hydrogen Evolution and Oxidation on Doped Transition Metal Sulfides. J. Electrochem. Soc. 2013, 160, A1902–A1906. Campbell, C. T.; Campbell, J. M.; Dalton, P. J.; Henn, F. C.; Rodriguez, J. A.; Seimanides, S. G. Probing Ensemble Effects in Surface Reactions. 1. Site-Size Requirements for the Dehydrogenation of Cyclic Hydrocarbons on Platinum(111) Revealed by Bismuth Site Blocking. J. Phys. Chem. 1989, 93, 806–814. Lessard, J. Electrocatalytic Hydrogenation. In Organic Electrochemistry; Hammerich, O., Speiser, B., Eds.; Boca Raton, 2015; pp 1657–1672. Markovic, N. M.; Ross Jr., P. N. Surface Science Studies of Model Fuel Cell Electrocatalysts. Surf. Sci. Rep. 2002, 45, 117– 229. Van Der Niet, M. J. T. C.; Den Dunnen, A.; Juurlink, L. B. F.; Koper, M. T. M. Co-Adsorption of O and H2O on Nanostructured Platinum Surfaces: Does OH Form at Steps? Angew. Chemie Int. Ed. 2010, 49, 6572–6575. Chang, S. C.; Weaver, M. J. In Situ Infrared Spectroscopy at Single-Crystal Metal Electrodes: An Emerging Link between Electrochemical and Ultrahigh-Vacuum Surface Science. J. Phys. Chem. 1991, 95, 5391–5400. Hansen, M. H.; Jin, C.; Thygesen, K. S.; Rossmeisl, J. Finite Bias Calculations to Model Interface Dipoles in Electrochemical Cells at the Atomic Scale. J. Phys. Chem. C 2016, 120, 13485–13491. Hansen, M. H.; Nilsson, A.; Rossmeisl, J. Modelling pH and Potential in Dynamic Structures of the Water/Pt(111) Interface on the Atomic Scale. Phys. Chem. Chem. Phys. 2017, 19, 23505– 23514. Kristoffersen, H. H.; Vegge, T.; Hansen, H. A. OH Formation and H2 Adsorption at the Liquid Water-Pt(111) Interface. Chem. Sci. 2018, 9, 6912–6921. Yeh, I. C.; Hummer, G. System-Size Dependence of Diffusion Coefficients and Viscosities from Molecular Dynamics Simulations with Periodic Boundary Conditions. J. Phys. Chem. B 2004, 108, 15873–15879. Yoon, Y.; Rousseau, R.; Weber, R. S.; Mei, D.; Lercher, J. A. First-Principles Study of Phenol Hydrogenation on Pt and Ni Catalysts in Aqueous Phase. J. Am. Chem. Soc. 2014, 136, 10287–10298. Silbaugh, T. L.; Campbell, C. T. Energies of Formation Reactions Measured for Adsorbates on Late Transition Metal Surfaces. J. Phys. Chem. C 2016, 120, 25161–25172. Honkela, M. L.; Björk, J.; Persson, M. Computational Study of the Adsorption and Dissociation of Phenol on Pt and Rh Surfaces. Phys. Chem. Chem. Phys. 2012, 14, 5849. Morin, C.; Simon, D.; Sautet, P. Chemisorption of Benzene on Pt(111), Pd(111), and Rh(111) Metal Surfaces : A Structural and Vibrational Comparison from First Principles. J. Phys. Chem. B 2004, 108, 5653–5665. Liu, W.; Carrasco, J.; Santra, B.; Michaelides, A.; Scheffler, M.; Tkatchenko, A. Benzene Adsorbed on Metals: Concerted Effect of Covalency and van Der Waals Bonding. Phys. Rev. B Condens. Matter Mater. Phys. 2012, 86, 1–6. Carrasco, J.; Liu, W.; Michaelides, A.; Tkatchenko, A. Insight

ACS Paragon Plus Environment

Page 9 of 9 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

(46)

(47) (48)

(49) (50)

(51) (52)

(53) (54)

into the Description of van Der Waals Forces for Benzene Adsorption on Transition Metal (111) Surfaces. J. Chem. Phys. 2014, 140, 084704. Song, Y.; Gutiérrez, O. Y.; Herranz, J.; Lercher, J. A. Aqueous Phase Electrocatalysis and Thermal Catalysis for the Hydrogenation of Phenol at Mild Conditions. Appl. Catal. B Environ. 2016, 182, 236–246. Chou, P.; Vannice, M. A. Benzene Hydrogenation over Supported and Unsupported Palladium. II. Reaction Model. J. Catal. 1987, 107, 140–153. Stoerzinger, K. A.; Favaro, M.; Ross, P. N.; Yano, J.; Liu, Z.; Hussain, Z.; Crumlin, E. J. Probing the Surface of Platinum during the Hydrogen Evolution Reaction in Alkaline Electrolyte. J. Phys. Chem. B 2018, 122, 864–870. McCrum, I. T.; Janik, M. J. First Principles Simulations of Cyclic Voltammograms on Stepped Pt(553) and Pt(533) Electrode Surfaces. ChemElectroChem 2016, 3, 1609–1617. Li, J.; Ghoshal, S.; Bates, M. K.; Miller, T. E.; Davies, V.; Stavitski, E.; Attenkofer, K.; Mukerjee, S.; Ma, Z. F.; Jia, Q. Experimental Proof of the Bifunctional Mechanism for the Hydrogen Oxidation in Alkaline Media. Angew. Chemie - Int. Ed. 2017, 56, 15594–15598. Rossmeisl, J.; Chan, K.; Skúlason, E.; Björketun, M. E.; Tripkovic, V. On the pH Dependence of Electrochemical Proton Transfer Barriers. Catal. Today 2016, 262, 36–40. Velasco-Velez, J.-J.; Wu, C. H.; Pascal, T. A.; Wan, L. F.; Guo, J.; Prendergast, D.; Salmeron, M. The Structure of Interfacial Water on Gold Electrodes Studied by X-Ray Absorption Spectroscopy. Science. 2014, 346, 831–834. Schultz, Z. D.; Shaw, S. K.; Gewirth, A. A. Potential Dependent Organization of Water at the Electrified Metal-Liquid Interface. J. Am. Chem. Soc. 2005, 127, 15916–15922. Christmann, K. Adsorbate Properties of Hydrogen on Solid Surfaces. In Landolt-Börnstein, Group III, Vol. 42, Physics of

(55) (56)

(57) (58)

(59) (60) (61) (62)

(63)

Covered Solid Surfaces, Subvolume A, Part 5(H.-P. Bonzel); 2006. Christmann, K. Interaction of Hydrogen with Solid Surfaces. Surf. Sci. Rep. 1988, 9, 1–163. Sanyal, U.; Song, Y.; Singh, N.; Fulton, J. L.; Herranz, J.; Jentys, A.; Gutiérrez, O. Y.; Lercher, J. A. Structure Sensitivity in Hydrogenation Reactions on Pt/C in Aqueous-Phase. ChemCatChem 2018, in press: DOI: 10.1002/cctc.201801344. Perdew, J.; Burke, K.; Ernzerhof, M. Generalized Gradient Approximation Made Simple. Phys. Rev. Lett. 1996, 77, 3865– 3868. Vandevondele, J.; Krack, M.; Mohamed, F.; Parrinello, M.; Chassaing, T.; Hutter, J. Quickstep: Fast and Accurate Density Functional Calculations Using a Mixed Gaussian and Plane Waves Approach. Comput. Phys. Commun. 2005, 167, 103–128. Lippert, G.; Hutter, J.; Parrinello, M. A Hybrid Gaussian and Plane Wave Density Functional Scheme. Mol. Phys. 1997, 92, 477–488. VandeVondele, J.; Hutter, J. Gaussian Basis Sets for Accurate Calculations on Molecular Systems in Gas and Condensed Phases. J. Chem. Phys. 2007, 127, 114105. Goedecker, S.; Teter, M.; Hutter, J. Separable Dual-Space Gaussian Pseudopotentials. Phys. Rev. B 1996, 54, 1703–1710. Grimme, S.; Antony, J.; Ehrlich, S.; Krieg, H. A Consistent and Accurate Ab Initio Parametrization of Density Functional Dispersion Correction (DFT-D) for the 94 Elements H-Pu. J. Chem. Phys. 2010, 132, 154104. Garcia-Araez, N. Enthalpic and Entropic Effects on Hydrogen and OH Adsorption on Pt(111), Pt(100), and Pt(110) Electrodes as Evaluated by Gibbs Thermodynamics. J. Phys. Chem. C 2011, 115, 501–510.

ACS Paragon Plus Environment