Fig. 1. Mechanistic insights into the direct H2O2 synthesis over Pd catalysts in aqueous phase. (a) A plausible sequence of elementary steps forming H2O2 and H2O, featuring ionization of H-adatoms. (b) Pd NPs acting as a microelectrode for the two half-reactions. (c) A schematic diagram of the free energy landscape for H2O2 and H2O formation on Pd NPs. Adapted with permission from Refs. [45,47]. Copyright American Chemical Society (2016, 2021).

Fig. 2. Initial TOF for aqueous-phase hydrogenation of various carbonyl compounds over monometallic catalysts at 373 K with an aqueous solution containing 5 wt% reactant as the feed. Plotted using the numerical data shown in Ref. [60].

Fig. 3. Influence of pH on the catalytic rates of benzyl alcohol hydrogenolysis and furfural hydrogenation at Pd-aqueous interfaces. (a) TOF as a function of pH in the catalytic conversion of benzyl alcohol to toluene on Pd/C. (b) Log(TOF) as a function of pH to obtain the apparent reaction order in bulk hydronium ion. (c) TOF of furfural hydrogenation as a function of buffer pH and the corresponding open circuit potentials at different H2 pressures. (d) Isotherms of furfural adsorption on Pd/C in 0.1 mol L-1 phosphate buffer solution at room temperature. Reproduced with permission from ref. [73,74]. Copyright retained by the authors of the cited articles under Creative Commons Licenses.

Fig. 4. Low-energy sputtering (LES) and reactive ion scattering (RIS) measurements of adsorbed H2O (1.2 monolayer) with or without co-adsorbed H or D (0.75 monolayer equivalent) on Pt(111). (a) LES and RIS spectra obtained on a Pt(111) surface with co-adsorbed H and H2O at ≤ 90 K (top), when heated at 140 K for 100 s and briefly at 150 K (middle), and only H2O was adsorbed on a bare Pt(111) surface and heated at 140 K for 100 s and briefly at 150 K (bottom). (b) LES and RIS spectra obtained from a Pt(111) surface with co-adsorbed D and H2O at < 90 K (upper) and for the sample heated at 140 K for 100 s and briefly at 150 K (lower). Adapted from Ref. [89,91]. Copyright 2015 Wiley-VCH.

Fig. 5. A quantitative analysis of H* ionization and interfacial proton formation on transition metal surfaces. (a) A three-dimensional thermochemical construct of homolytic/heterolytic dissociation of HM (i.e., H* adatom on the metal surface). (b) Periodic trends of pKa values (in bracket), which is related to the free energy for creating a H+sol on closed-packed Group VIII transition metal surfaces. (c) Computed pKa values of H-adatom as a function of H* coverage. (d) DFT-derived work functions of closed-packed transition metal surfaces, phenol covered Ru(001) surface, and phenoxy covered Ru(001) surface, represented by different symbols. Reproduced with permission from ref. [78]. Copyright 2022 Elsevier.

Fig. 6. Water promotion of vapor-phase hydrogenation of butanal over Ru/SiO2. (a,b) Dependence of hydrogenation turnover rates and the corresponding promotional ratios, γ, to (a) changes in the H2O pressure and (b) with different combinations of hydrogen (H2 or D2) and water (H2O and D2O). (c) Illustrates the differences in the alkoxy pathway with and without water vapor. Adapted with permission from Ref. [92]. Copyright 2021 Elsevier.

Fig. 7. Demonstrating the PCET mechanism of aqueous-phase 4-nitrophenol hydrogenation over Pd/C using unbiased electrochemical cells. (a) Schematic illustration of H-cell and S-cell setups. (b) Comparison of the reaction orders in hydrogen, proton and 4-nitrophenol using Pd/C (5 wt%)-carbon cloth electrode combination in the H-cell and Pd/C (0.075 wt%) in the S-cell. (c) Comparison of the hydrogenation rates of 4-nitrophenol in H- and S-cells using different electrode combinations in the pH range 0-2 and the relationship between electron-transfer rates (based on the current between the anode and cathode) and hydrogenation rates in the H-cell. (d) A proposed 4-nitrophenol hydrogenation pathway driven by electron transfer. Adapted with permission from Ref. [93]. Copyright 2022 American Chemical Society.

Fig. 8. Aqueous-phase cinnamaldehyde hydrogenation over CNT-supported PtFe catalysts: (left panel) diagram of product distributions in the (control) experiments of reagent-D2O exchange with/without Pt3Fe/CNT and the catalytic hydrogenation reaction in D2O; (right panel) an illustration of water-involved hydrogen exchange and DFT-computed energetic profiles for C=C and C=O hydrogenation routes over the surface model of PtFe NPs with/without the assistance of water. Adapted with permission from Ref. [57]. Copyright 2018 Elsevier.

Fig. 9. Isotopic evidence for the Langmuir-Hinshelwood pathway for furfural hydrogenation over Pd/C (30?mmol L-1 furfural in 0.1? mol L-1 phosphate buffer solution with 10 mg?Pd/C at 1?bar H2 or D2 and room temperature). (a) Turnover frequency (TOF) of furfural hydrogenation in D2O/H2O and D2/H2 and (b-d) the mass spectra of the produced furfuryl alcohol (FOL). Mass spectra of (b) standard FOL dissolved in H2O and D2O in comparison with those of FOL produced form furfural in H2; and mass spectra of FOL produced from furfural at (c) pH 1.6 and (d) pH 5.8. Reproduced with permission from Ref. [73]. Copyright retained by the authors of the cited article under Creative Commons Licenses.

Fig. 10. Isotopic experiments used to elucidate the mechanism of liquid-phase benzaldehyde hydrogenation on Pd/C at 298 K. (a) The reaction pathway leading to the formation of D-benzaldehyde (C6H5CDO) in the presence of D2. (b) Formation rate of D-benzaldehyde and D-benzyl alcohol in either dioxane or D2O solvent under D2. (c) Kinetic isotope effect of H2O/D2O and H2/D2 on TOF of benzaldehyde hydrogenation. Adapted with permission from Ref. [94]. Copyright 2021 Springer Nature.

Fig. 11. DFT assessments of water effects on acetone hydrogenation on metallic surfaces. (a) Energetic span (in eV) for the acetone hydrogenation at water-free conditions (in black) and in the presence of one chemisorbed water (in grey) for a series of transition metal surfaces. Adding 11 water molecules or using an implicit continuum model did not lead to a significant change in the energetic span. (b) Variation of the energetic span of the alkoxy route upon water assistance as a function of the d-band center energy (Cu notably deviates from the linear correlation since its d-band is completely filled and as a consequence, the d-band model is not applicable). Adapted with permission from Ref. [61]. Copyright 2014 Royal Society of Chemistry.

Fig. 12. Similarities in the postulated pathways for the reductive hydrolysis of aryl ethers (R = phenyl, cyclohexyl, phenylethyl, and n-butyl) in water (a) and anisole in H2O/CH2Cl2 (b), both catalyzed over Pd surfaces (Pd/C catalysts). Adapted with permissions from Refs. [102,106]. Copyright 2017 Wiley VCH and Springer Nature.

Fig. 13. Hydrogen binding energy (HBE) on metal surfaces as an explanation for pH effects in hydrogenation reactions. (a) Initial TOFs of phenol hydrogenation over 5 wt% Pt/C at 353 K and 20 bar H2 and DFT-computed HBEs for Pt(110) plotted as a function of pH. (b) Replotted pH dependence as ln(TOF) versus HBE/RT. (c) A schematic enthalpy diagram of a hydrogenation step at different pH values. Adapted with permission from Ref. [72]. Copyright 2019, American Chemical Society.

Fig. 14. A detailed analysis of aqueous-phase furfural hydrogenation on Pd/C. (a) Effects of pH on K1o (the equilibrium constant for the first H-addition to adsorbed furfural), k2 (the rate constant for the rate-determining step, which is the second H-addition to a monohydrogenated furfural adspecies), and coverages (θFH: coverage of the monohydrogenated furfural species, θH: H* coverage) in furfural hydrogenation. (b) Correlation between the regressed K1o and KH2o (the equilibrium constant for the dissociative adsorption of H2). (c) Energy diagram based on experimentally determined values for the illustration of pH effects on hydrogenation kinetics. Adapted with permission from Ref. [73]. Copyright retained by the authors of the cited article under Creative Commons Licenses.

Fig. 15. Phenomena and origins of Br?nsted-acid-site (BAS)-promoted aqueous-phase hydrogenation of phenol over Pt/silicalite-1 and Pt/H-MFI catalysts. (a) TOF as a function of BAS concentrations. (b) Experimentally measured reaction energy profile of phenol hydrogenation on the non-acidic Pt/silicalite-1 via conventional hydrogenation pathway and on the most acidic Pt/HMFI-24 via a PCET-type pathway. (c) DFT-calculated electronic reaction energy profile for H-addition to the bound phenol via a PCET-type mechanism in the “11w-acidic” (11 water, 1 proton) zeolite and “8w-acidic” (8 water, 1 proton) zeolite models with the transition states for the 1st and 2nd H-addition via conventional H attack marked on the plot. Adapted with permission from Ref. [116]. Copyright 2021, Elsevier.

Fig. 16. DFT-calculated activation energies for the key elementary steps involved in gas-phase ($\Delta E_{\text {act }}^{\text {vac }}$, black bars) and aqueous-phase ($\Delta E_{a c t}^{a q}$, white bars) catalytic dehydrogenations of COH* and CH3OH*, as well as those calculated by subtracting the interaction energy of the reactants from that of the transition state (ΔΔEint, gray bars), on Pt(111) surface ($\Delta E_{a c t}^{a q}=\Delta E_{a c t}^{v a c}+\Delta \Delta E_{\text {int }}$), along with the corresponding transition states. Adapted with permission from Ref. [128]. Copyright 2019, Royal Society of Chemistry.

Scheme 1. Mechanistic ambiguities associated with hydrogenation and hydrogenolysis reactions at metal-aqueous interfaces, highlighting the classical LH and the PCET-type (electrochemical) mechanisms. X=Y represents a molecule with an unsaturated bond to be hydrogenated, while X-Y represents a molecule (bond) to be hydrogenolytically cleaved. Note that competitive adsorption of hydrogen and the substrate is not implied and the steps drawn in this scheme are not necessarily elementary steps.

Fig. 17. CO consumption rates (a) and selectivity to hydrocarbons (b,c) as a function of water volume fraction for Ru/CNT and Ru/CNT-Ox (CNT-Ox: surface-oxidized CNT). Reactions were performed in a batch reactor at 493 K, 800 psi H2/CO (4/1), 300 r min-1, conversion ～20%. Adapted with permission from Ref. [141]. Copyright 2020, American Chemical Society.

Fig. 18. Remarkable effects of water (as co-fed moisture) on CO oxidation rates over various Au-based catalysts reported in the literature: (a) the work of Haruta and coworkers [150]; (b) the work of Chandler and coworkers [162]; (c) the work of Iglesia and coworkers (the inset shows the reversible nature of water effects) [163]. Adapted with permission from Refs. [150,162,163]. Copyright Wiley VCH (2004), Elsevier (2012) and American Chemical Society (2018).

Fig. 19. H2O/D2O isotope effects on CO oxidation turnover rates measured on Au-based catalysts. (a) KIE over Au/Al2O3 (288 K; a feed containing 5 kPa CO, 2 kPa O2 with or without 0.5 kPa H2O/D2O [163]. (b) Average KIE with H2O- and D2O-exchanged Au/Al2O3 catalysts and effects of adding and removing 700 Pa H2O/D2O (293 K; a 1% CO and 20% O2 feed in N2 balance) [162]. Adapted with permission from Refs. [162,163]. Copyright 2012 Elsevier and 2018 American Chemical Society.

Fig. 20. A schematic illustration of the water adsorption on supported AuNPs for different moisture levels (I-IV) and qualitative dependence of the catalytic activity of Au catalysts on the relative humidity in the feed gas. Reprinted with permission from Ref. [144]. Copyright 2017, American Chemical Society.

Fig. 21. Turnover frequencies for the production of CO2 (a) and H2O2 (b) in the aqueous-phase CO oxidation over three Au catalysts at different pH values. Plotted using tabulated data in Ref. [174]. Note that CO2 was measured by the formation of Na2CO3 at pH = 14 and the TOF for H2O2 formation at pH = 0.3 on Au/Calgon and Au/carbon-WGC (World Gold Council) was shown as the upper limit (< 0.001 s-1) as both were below the detection limit.

Fig. 22. Selective glycerol oxidation to DHA catalyzed by supported Au catalysts. (a) A comparison of TOF over Au catalysts. (b) The variation of DHA selectivity (C%) as a function of glycerol conversion (note: the conversion was changed by varying the reaction temperature from 303 to 373 K). Plotted using data tabulated in Ref. [196].

Fig. 23. Promoting effects of water on the conversion (blue bars) and benzaldehyde selectivity (orange bars) over Au/TiO2 for the selective oxidation of benzyl alcohol in pure water, pure p-xylene and biphasic mixtures of water and p-xylene. Plotted using data tabulated in Ref. [185].

Fig. 24. Proposed general schemes of alcohol oxidation over polymer-protected Au NP catalysts (a,b) (a: PVP-stabilized, b: PS-stabilized) [215] and PVP-stabilized atomically precise Au24 cluster (c) [214], and a DFT-derived complete mechanism for the aerobic oxidation of 1-phenylethanol over a Au252 cluster model (d) [219]. Reprinted with permissions from Refs. [214,215,219]. Copyright American Chemical Society (2005, 2011 and 2021).

Fig. 25. Depictions of the reaction mechanisms for CH3OH oxidation on Au/TiO2/O2 in the gas phase (a), in neutral water (b), and in basic water (c), as extracted from the AIMD simulations. Adapted with permission from Ref. [225]. Copyright 2018, Wiley-VCH.

Fig. 26. Aerobic alcohol oxidation in biphasic water-organic mixtures. (a) A schematic model of a multiphase reaction system for the selective oxidation of benzyl alcohol to benzaldehyde over Au/TiO2 [185]. (b) Postulated phase transfer catalysis for the ‘switch’ effect of water for the catalytic activity of Rh NPs in the aerobic oxidation of 1-phenylethanol to acetophenone in a biphasic toluene-water mixture (the two solvent components are nearly immiscible) [208], which involves the alcohol transfer to the water phase (Step 1), the oxidation reaction over Rh surface in water (Step 2) and the extraction of the largely immiscible acetophenone product in water by toluene (Step 3). Reprinted with permissions from Refs. [185] and [208]. Copyright 2008 Elsevier and 2017 Royal Society of Chemistry.

Fig. 27. Alcohol oxidation at solid-gas and solid-liquid interfaces with mesoporous silica-supported Pt NPs as the catalyst. (a) Effects of H2O addition on the TOF of gas-phase and liquid-phase oxidation of 2-propanol (left) and 1-propanol (right). (b) Apparent activation energy for 1-propanol oxidation on 4.5 nm Pt in gas and liquid phases with/without H2O. (c) Dependence of apparent activation energy on the average size of Pt NPs in neat 2-propanol and 2-propanol/water and 2-propanol/n-heptane mixtures. Reproduced or replotted with permission from Refs. [232] and [233]. Copyright American Chemical Society (2014, 2019).

Fig. 28. Temperature-programmed surface reactions of methane oxidation on Cu-CHA with isotopically labeled reactants (98% methane, 2% water, 400 ± 50 ×10-6 bar dioxygen for all mixtures). (a) CH4-D2O-O2. (b) CD4-H2O-O2. (c) CH4-H218O-O2. (d) CH4-H2O-18O2. Adapted with permission from Ref. [247]. Copyright 2021, Elsevier.

Fig. 29. Comparison of the methanol selectivity-methane conversion relationship at 323 K for gas- (blue) and aqueous- (cyan) phase reactions using DFT-computed free energy of activation ΔGDFT?, with a correction on the cyan line equal to the solvation free energy of methanol at 298 K (-0.22 eV). Only the experimental data from reactions conducted in aqueous conditions are plotted. Reprinted with permission from Ref. [268]. Copyright 2018, American Chemical Society.

Scheme 2. Mechanistic ambiguities related to the activation of O2 and the roles of roles of O2-derived species in activating the organic substrates in catalytic aerobic oxidation at metal-aqueous interfaces (exemplified for alkanes and alkanols). Note that not all possible routes for O2 activation and aerobic oxidation are shown in this scheme.

Fig. 30. Cooperative redox enhancement effects on the catalytic activity of aerobic oxidation of oxygenates at metal-aqueous interfaces under alkaline conditions. (a) HMF conversion as a function of time over various catalyst formulations containing Au and Pd. (b) Reaction rates for the aqueous-phase aerobic oxidation of various substrates over mono- and bimetallic formulations. (c) A proposed reaction scheme illustrating the separated dehydrogenation and oxygen reduction reactions over a physical mixture of Au/C and Pd/C. (d) A correlation between the thermo- and electro-catalytic HMF oxidation rates. Reprinted with permission from Ref. [269]. Copyright 2022, Springer Nature.

Fig. 31. Mechanistic connections between thermo- and electrocatalytic aerobic oxidation of different substrates in water. (a) A short-circuit (SC) model viewing a thermocatalytic oxidation as two coupled electrochemical half-reactions (Ecat refers to the potential of the catalyst during thermochemical turnover). (b) Independent electrical polarization curves for each half-reaction predict the mixed potential and the reaction rate during thermochemical catalysis (ELSV refers to the catalyst potential predicted by applying the current-matching condition to the electrochemical polarization curves). (c,d) The SC model predicts mixed potentials and rates of oxidation catalysis, showing a marked similarity between ELSV and the measured Ecat, and between the electrocatalytic TOF at Ecat and the thermocatalytic TOF of various aerobic oxidation reactions. Reprinted with permission from Ref. [23]. Copyright 2021, Springer Nature.

Fig. 32. (a) AIMD simulations of Pd(111) surface in the presence of water and bulk and surface hydrogen at 473 K. (b) Simulated Pd-O radial distribution function between surface Pd atoms and oxygen in water, with the arrows pointing to the first shell distance showing the increased hydrophobicity as the surface concentration of hydrogen increases. (c) Fraction of Ni0 formed as a function of time for selected Ni catalysts under 50 bar H2 at 473 K and in 5 wt% phenol/H2O under 35 bar H2 at 473 K from in situ EXAFS analysis (note: at zero time the reactor was at ambient temperature and the reaction temperature was reached in ca. 12 min). Adapted with permissions from Refs. [279,283]. Copyright American Chemical Society (2013) and Wiley-VCH (2015).

Fig. 33. Snapshots from the AIMD simulations showing the initial (a), intermediate (b), and final (c) configurations for water dissociation at the periphery of a Pt6 cluster supported by CeO2. Adapted with permission from ref. [295]. Copyright 2016, American Chemical Society.

Scheme 3. The overall Born-Haber thermodynamic cycle for the free energy required to create hydronium ions from adsorbed hydrogen (drawn as one adsorbed on a three-fold hollow site) in the presence of water or any solvent in general, with the shaded box representing the water/solvent layer. Note: the only steps (in boldfaced texts and figures) which change on moving to different metals are the desorption of H? and the work function of the metal. While the desorption of H? changes with the metal substrate used, the change is small compared with the changes in the work function of the metal. Gas-phase ionization and H+ solvation (a value shown for H+ solvation in water) are independent of the metal identity (italic texts and figures).

Fig. 34. Water-promoted partial oxidation of methane over atomically dispersed Au1/BP catalyst under light irradiation. (a) Comparison of methanol yields under different solvents and reactants at 363 K for 2?h. (b) In situ ESR spectra of BP and Au1/BP nanosheets under different conditions in the presence of 5,5-dimethyl-1-pyrroline-N-oxide (DMPO), which was selected for the detection of superoxide ions (O2-) and hydroxyl radicals (?OH). (c,d) In situ DRIFTS spectra of Au1/BP nanosheets purged by 1?bar of different gases at 363 K in the dark under light irradiation. Adapted with permission from Ref. [298]. Copyright retained by the authors of the cited article under Creative Commons Licenses.

Fig. 35. A Born-Haber thermochemical cycle linking gas-phase thermodynamic quantities with liquid-phase thermodynamic quantities for quasi-equilibrated steps in catalytic hydrogenation of carbonyl compounds at the interface of protic solvent (water and alcohols) and Ru metal NPs. Reprinted with permission from Ref. [77]. Copyright 2019, American Chemical Society.

Fig. 36. Charge analyses and transfer during interfacial water dissociations. (a) Bader charge analysis in aqueous phase. (b) The linear relationship between the charge and the number of dissociated interfacial waters. (c) The schematic charge transfer caused by spatial separation of Hδ+ and OHδ- during the nth interfacial water dissociation. Adapted with permission from Ref. [328]. Copyright 2021, Elsevier.