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A Simple Liquid−Liquid Biphasic System for Hydrogen Peroxide Generation Wojciech Adamiak,*,† Justyna Jedraszko,† Wojciech Nogala,† Martin Jönsson-Niedziolka,† Saustin Dongmo,‡ Gunther Wittstock,‡ Hubert H. Girault,§ and Marcin Opallo† †

Institute of Physical Chemistry, Polish Academy of Sciences, Kasprzaka 44/52, 01-224, Warsaw, Poland Institute of Chemistry, Center of Interface Science, Faculty of Mathematics and Natural Science, Carl von Ossietzky University of Oldenburg, D-26111 Oldenburg, Germany § LEPA, Ecole Polytechnique Federale de Lausanne, CH B2 401 (Bâtiment CH) Station 6 CH-1015 Lausanne, Switzerland Downloaded by SUNY UPSTATE MEDICAL UNIV on September 7, 2015 | http://pubs.acs.org Publication Date (Web): August 19, 2015 | doi: 10.1021/acs.jpcc.5b06620

‡

S Supporting Information *

ABSTRACT: Although electron transfer reactions at liquid−liquid interfaces have been thoroughly studied in the presence of deliberately added potential determining ions or phase transfer catalysts, little is known about these reactions in the absence of the above species. Here, we report the formation of hydrogen peroxide (H2O2) in a liquid−liquid two-phase system with only an electron donor (decamethylferrocene) solution in trifluorotoluene (TFT) and a proton donor solution (HClO4) in water. To detect H2O2, we used fluorescent microscopy and scanning electrochemical microscopy (SECM). We applied a potential sweep program to the SECM probe to overcome the electrode deactivation. To provide insight into the reaction rate and mechanism, we fitted SECM results to finite elements simulations. From the results of UV−vis spectroscopy, we determined a H2O2 partition coefficient of 0.03 and the standard Gibbs energy of H2O2 transfer from water to TFT as 9.5 kJ/mol. The most important conclusion from this work is that the studied system provides conditions for spontaneous transfer of protons from the aqueous to the organic phase, even in the absence of deliberately added potential determining ions or phase transfer catalysts.

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INTRODUCTION Hydrogen peroxide is produced on a large scale in an anthraquinone oxidation process, which involves catalytic hydrogenation and oxidation of the alkylanthraquinone precursor dissolved in an organic solvent mixture, followed by the extraction of H2O2 from the organic to the aqueous phase.1,2 The process is a multistep method that requires a significant energy input and generates waste, which have a negative effect on the overall production cost, as has been reviewed by Campos-Martin et al.3 A promising alternative to the anthraquinone process can be the generation of hydrogen peroxide in oil−water two-phase systems where the oil phase contains an electron donor, for example, decamethylferrocene (DMFc), and the water phase contains acid.4,5 In the presence of oxygen in such systems, the following reaction takes place

been used as water immiscible molecular solvents for H2O2 generation. The latter seems to be more advantageous than the former due to its lower vapor pressure,6,7 which allows the generation of a more stable liquid−liquid interface.5 In terms of liquid interface stability, the use of nonvolatile room-temperature ionic liquids appears to be an attractive alternative to molecular solvents. However, our recent studies have shown that care must be taken when choosing the ionic liquids, because their constituent ions can be transferred to the aqueous phase to maintain electroneutrality of the organic phase during the two-phase reaction.8 In the long term, such transfer can lead to significant changes in the compositions of both phases and would result in liquid interface instabilities. In the case of molecular solvents, their molecules are electrically neutral, and so they cannot be transferred electrochemically during the reaction. Depending on the ionic composition of the aqueous and organic phases, hydrogen peroxide formation occurs in the bulk organic phase as a result of a homogeneous electron transfer reaction5 or as a result of an heterogeneous electron transfer reaction at the liquid−liquid interface.4 The first mechanism requires protons to be transferred from the aqueous to the

2DMFc(oil) + 2H+(water) + 2X−(water) + O2 (oil) → 2DMFc+(oil) + H 2O2 (water) + 2X−(oil)

(1)

and the oil−water interface acts as a defect-free liquid membrane that allows separation of the reagents. The advantage of the two-phase method over the anthraquinone process is that it is a single-step process that does not require a noble metal catalyst and hydrogen gas. Also, it does not require large facilities with huge capital costs as it can be performed on a smaller scale as an on-site production. So far, 1,2dichloroethane (DCE)4 or α,α,α-trifluorotoluene (TFT)5 has © 2015 American Chemical Society

Received: July 10, 2015 Revised: August 7, 2015 Published: August 10, 2015 20011

DOI: 10.1021/acs.jpcc.5b06620 J. Phys. Chem. C 2015, 119, 20011−20015

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The Journal of Physical Chemistry C

diameter with filament from Harvard Apparatus) using a 5 mm pulling distance. To avoid penetration of the aqueous phase into the pipet, and to obtain a stable TFT−water (W) interface, the inner walls of the pipet were hydrophobized by silanization with methyltrimethoxysilane vapors.11 The optical micrographs were obtained by a DMIRE2 microscope in an inverted configuration (Leica Microsystems GmbH). Samples were excited with a tungsten lamp with a dichromic filter set for excitation with the wavelength range of 500−600 nm. The detection was done for wavelengths from 575−650 nm by a DC152QC-FI sCMOS camera (Andor Technology) attached to the microscope. The objective was an HC PL FLUOTAR (Leica, 10× magnification, numerical aperture NA = 0.3). The recorded intensities were converted to false color images. All measurements were performed at room temperature (23 ± 2 °C).

organic phase, which, in turn, can be achieved by addition of a phase transfer catalyst, like lithium tetrakis(pentafluorophenyl)borate (LiTB), to the aqueous phase.5 The protonation of TB− anions is responsible for proton transfer to the organic phase in the form of HTB acid. Alternatively, one can polarize the liquid interface by the use of tetrakis(pentafluorophenyl)borate containing salts in both liquid phases.5 The established Galvani potential difference between organic and aqueous phase is lower than the transfer potential of protons, which ensures favorable conditions for their spontaneous transfer from the aqueous to the organic phase. On the other hand, when TB− anions are replaced by ClO4−, the Galvani potential difference is higher than the transfer potential of protons, which prevents them from transferring to the organic phase.4 In such a case, reaction 1 occurs at the liquid interface as a heterogeneous electron transfer. The only two-phase system with either a DCE−water or a TFT−water interface where hydrogen peroxide formation has not been investigated so far is the one with no deliberately added potential determining ions or phase transfer catalysts. In such a case, the liquid−liquid interface is likely to be polarized chemically due to partial dissolution of the proton donor between two liquid phases.9 The aim of this work was to determine whether such conditions are still sufficient to obtain H2O2 in the two-phase reaction and to fill the gap in fundamental knowledge about electron transfer reactions at soft interfaces. In fact, the two-phase system discussed in this work is also the simplest one for electron transfer reactions at soft interfaces for H2O2 generation, as it involves only the reagents and a pair of solvents. This property can be especially attractive from an industrial application point of view.

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RESULTS AND DISCUSSION Chemical Detection. At first, we performed a two-phase reaction with a DMFc solution in TFT and an aerated aqueous solution of HClO4. TFTa water-immiscible solventis less toxic than the previously used DCE.4 After 24 h, the color of the organic phase changed from yellow (Figure 1, flask 1A) to

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EXPERIMENTAL SECTION Chemicals. Decamethylferrocene (DMFc, 97%) was purchased from ABCR, and α,α,α-trifluorotoluene (TFT, ≥99%) was from Acros Organics. 70% HClO4, NaClO4, tetrahexylammonium chloride (THxACl), and horseradish peroxidase were from Sigma-Aldrich. Amplex UltraRed fluorescent dye was obtained from Life Technologies, and KI was from ChemPur. All chemicals were used as received without further purification. Tetrahexylammonium perchlorate (THxAClO4) was prepared by metathesis of THxACl and HClO4 and recrystallized twice from a mixture of ethyl acetate and ethanol.10 All aqueous solutions were prepared with demineralized water from an ELIX system (Millipore). Apparatus and Procedures. Electrochemical Measurements. SECM measurements were carried out with a CHI900B SECM workstation (CH Instruments). Pt microelectrodes for SECM experiments were made by sealing a Pt wire (25 μm diameter, Goodfellow, England) using a PC-10 micropipet puller (Narishige) in borosilicate glass capillaries and subsequent polishing. A Pt wire and Hg|Hg2SO4|K2SO4(sat) were used as counter and reference electrodes, respectively, and were immersed in the aqueous phase. Hg|Hg2SO4| K2SO4(sat) was chosen instead of Ag|AgCl|3 M KCl to avoid possible contribution of Cl− oxidation to the measured current. The Pt microelectrode SECM tip was inserted into the aqueous phase, and its position was controlled by stepper motors in the X, Y, and Z directions. Fluorescence Detection. The experiments with Amplex UltraRed were performed with glass pipets (orifice diameter ca. 200 μm) prepared with a PC-10 micropipet puller (Narishige) from borosilicate glass (1.17 mm inner diameter, 1.5 mm outer

Figure 1. Photographs of shake-flasks experiments in TFT−W biphasic systems. Flask 1A contains a TFT solution of 5 mM DMFc (bottom phase) in contact with aerated 0.1 M HClO4 aqueous solution (upper phase) right after the two solutions were brought in contact. Flask 1B is flask 1A after 24 h of the two-phase reaction. Flask 1C contains the aqueous phase collected from flask 1B with an addition of 100 μL of 0.1 M KI and 10% starch. Flasks in row 2 are the same as flasks in row 1, except that the aqueous phase contains aqueous 0.1 M NaClO4.

green (Figure 1, flask 1B), indicating that DMFc was oxidized. In contact with pure water, the organic phase remained yellow after 24 h (Figure 1, flasks 2A, 2B). This observation suggests that DMFc requires not only molecular O2 as oxidizer but also a proton donor dissolved in the aqueous phase. When the aqueous phase was separated from the organic phase after the two-phase reaction, and KI and starch were added, the solution color changed to purple (Figure 1, flask 1C). This is due to formation of a complex between the starch and I3− which is formed in the specific reaction sensitive to H2O2: 3I− + H 2O2 + 2H+ → I3− + 2H 2O 20012

(2)

DOI: 10.1021/acs.jpcc.5b06620 J. Phys. Chem. C 2015, 119, 20011−20015

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The Journal of Physical Chemistry C

Figure 2. (A) Amperometric approach curves to TFT−W interface recorded 30 s after the interface formation. The red curve was recorded on the activated Pt electrode, whereas the blue curve was recorded on the nonactivated electrode (see the text for details). (B) Comparison between experimental approach curve from (A) (red curve) and simulated approach curves to a 50 μm thick uniform reaction zone model interface (green triangles) and to a geometrically flat H2O2 source (black triangles).

To determine the H2O2 flux and provide additional insight into the reaction mechanism, the experimental approach curve was simulated using COMSOL Multiphysics software (see the Supporting Information for simulation details). Since the H2O2 generation takes place at (or close to) the liquid−liquid interface, a simplified model taking into account only the aqueous solution was used first. In the simplest version of the simulations, the H2O2 generation was modeled as a constant flux from the liquid−liquid interface. As can be seen in Figure 2B (black triangles), the current decreases dramatically close to the interface, which can be assigned to negative feedback with the SECM tip hindering diffusion of analyte.16 Such a behavior is not seen in the experimental curve, which means that H2O2 diffuses to the tip through the liquid−liquid interface from the organic phase. To simulate such a behavior, we assigned a thin reaction zone below the liquid−liquid interface where the H2O2 is generated. In this zone, a constant H2O2 volume source was assumed. In this model, a much smaller current drop for shortest distances was seen (green triangles in Figure 2B). Since only 3% of H2O2 partitions into the organic phase (see the Supporting Information), the total H2O2 generation rate is higher than estimated from the measured flux. The value of H2O2 flux obtained from simulations is 3 × 10−12 mol·cm−2·s−1. At distances larger than 800 μm, the experimental curve shows significantly higher anodic current than the calculated ones. This experimental background current is probably due to the aforementioned oxidation of the Pt surface during chronoamperometry.12 Close to the interface, the experimental tip current significantly exceeds simulated values. This is probably caused by a positive feedback, because H2O2 oxidation at the tip produces additional oxygen flux (eq 3), which is a substrate for H2O2 generation at the liquid−liquid interface. However, the reaction is kinetically sluggish and far from being limited by the supply of oxygen. We did perform simulations where the H2O2 generation rate was dependent on the oxygen concentration assuming pseudo-first-order reaction (O2 → H2O2, since H+ is available in large excess), Even with additional flux of O2 generated at the SECM tip by H2O2 oxidation, no significant difference was seen between this model and the one where the H2O2 generation rate was set independent of O2 generated at the tip (not shown). The

No color change of the aqueous phase is observed in the absence of H+ (Figure 1, flask 2C), indicating that H2O2 is not formed under these conditions. SECM Detection. Quantitative detection of H2O2 was performed amperometrically by SECM with a Pt tip positioned above the TFT−W interface. At sufficiently positive potential, oxidation of H2O2 was expected to occur at the tip: H 2O2 → O2 + 2H+ + 2e−

(3)

The main drawback of the constant-potential amperometric detection of H2O2 is that, at positive potentials, the surface of the Pt tip is deactivated due to formation of platinum oxides.12 Improved analytical performance of the Pt electrode has been achieved either by deposition of mesoporous Pt on the Pt surface13 or by controlled anodic dissolution of the electrode material.14 In our experiments, however, we did not change the geometry of the Pt surface itself. Instead, we have used a systematic cleaning of the electrode surface before recording the current at a given distance in a similar manner as that in ref 15. This procedure consists of the following steps: 1. A potential scan from +0.6 to −0.6 V at 500 mV·s−1 at a given distance, to reduce the Pt oxide layer and obtain an active Pt surface (Figure S1A). 2. Immediate chronoamperommetry on the freshly activated Pt surface is made at +0.6 V (Figure S1B). A diffusion-limited anodic current corresponding to H2O2 oxidation (eq 3) is measured, and the value at time = 5 s is taken to plot the approach curve. 3. 10 μm tip movement toward the liquid−liquid interface, followed by repetitions of the above steps until the entire approach curve is recorded. A plot of the tip current as a function of the approximate tipto-interface distance is presented as the red curve in Figure 2. The anodic tip current increases as the tip approaches the TFT−W interface, indicating H2O2 evolution at the interface. This current is ca. 5 times higher than the current recorded on the nonactivated electrode kept at constant potential (Figure 2A, blue curve). This clearly shows that the applied methodology is more advantageous than a traditional amperometric approach without tip activation. 20013

DOI: 10.1021/acs.jpcc.5b06620 J. Phys. Chem. C 2015, 119, 20011−20015

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instead coupled to ClO4− transfer across the liquid−liquid interface (eq 1). The electron transfer reaction can proceed homogeneously in the organic phase or heterogeneously at the liquid interface. In the first case, the electron transfer reaction requires H+ to be present in the organic phase. Initially, H+ ions are present only in the aqueous phase, but they can be transferred to the organic phase by HClO4 partitioning. Since the volumes of the organic and aqueous phases are equal, one can calculate concentration of H+ in the organic phase from the following equation

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apparent positive feedback on the experimental curve is, therefore, likely caused by a deformation of the interface pushed by the tip and/or formation of a thin layer of the aqueous phase at its surface upon TFT phase penetration.17 Fluorescent Detection. Formation of H2O2 in the twophase reaction was also confirmed by fluorescent microscopy using the specific reaction of Amplex UltraRed dye oxidation by H2O2, catalyzed by horseradish peroxidase.8,18 In the experiment, a filled glass capillary with TFT solution containing DMFc was immersed into an aqueous solution containing HClO4, horseradish peroxidase, and the dye. Strong fluorescence was observed at the end of the capillary a few seconds after the pipet was immersed into the aqueous phase (Figure 3A), indicating H2O2 formation.8 No fluorescence was

c H+(TFT) = c H0+(W) − c H+(W)

(4) +

where cH+(TFT) is the equilibrium concentration of H in TFT, and cH0 +(W) and cH+(W) are the initial and equilibrium concentrations of H+ in W, respectively. Assuming equal activity coefficients in both phases, the equilibrium concentration of H+ in W is given by9 c H+(W) =

c H0+(W) 1 + e H+

(5)

where ⎡⎛ F ⎞ TFT ⊖ ⎤ ⎟(Δ ϕH+ − ΔTFT ϕ)⎥ e H+ = exp⎢⎜ W W ⎣⎝ RT ⎠ ⎦

(6)

TFT In the above equation, ΔW ϕ is the Galvani potential ⊖ difference between TFT and W, and ΔTFT W ϕH+ is the standard + TFT transfer potential of H . The value of ΔW ϕ can be calculated by numerically solving the equation9

∑ i=1

⎡ 1 + exp⎣

zici0(W) ziF RT

⎤ ⊖ TFT (ΔTFT W ϕi − ΔW ϕ)⎦

( )

=0 (7)

where subscript i denotes all ionic species present in the given two-phase system (in our case, H+ and ClO4−). By taking TFT ⊖ ⊖ 5 5 ΔTFT W ϕH+ = −0.523 V and ΔW ϕClO4− = −0.235 V, one obtains ΔTFT W ϕ = −0.144 V. Next, from eqs 4−6, one calculates the concentration of H+ in TFT before the reaction, cH+(TFT) = 30 nM. As the reaction proceeds, the aqueous phase acts as a reservoir that continuously counteracts the consumption of H+. To determine the driving force of the homogeneous electron transfer reaction, one can calculate the standard redox potential TFT of oxygen reduction in TFT, [E⊖ O2/H2O2]SHE , from

Figure 3. (A) Fluorescence inside a micropipet filled with 5 mM DMFc in TFT immersed in the aqueous solution containing 0.1 M HClO4 and AmplexUltraRed fluorescent dye. (B) Fluorescence after deliberate addition of H2O2 to the bulk aqueous phase. The outline of the pipet is indicated by the white lines.

⊖ TFT W [EO⊖2 /H2O2]SHE = [EO⊖2 /H2O2]SHE − (ΔTFT W G H 2O2

observed in the absence of DMFc in TFT. The fluorescence is observed only in the organic phase at the end of the capillary and is likely to be due to partitioning of the dye into the TFT phase, as in a similar experiment presented in ref 8. This is reasonable considering that fluorescence is emitted only by the deprotonated form of the dye, which is not likely to exist in the strongly acidic aqueous phase. This hypothesis was verified in an experiment where H2O2 was deliberately added to the bulk aqueous phase. Again, fluorescence was observed only in the organic phase, indicating transfer of the deprotonated dye to the organic phase (Figure 3B). Reaction Mechanism. To maintain electroneutrality of the organic phase, continuous oxidation of DMFc to DMFc+ has to be coupled to either transfer of DMFc+ to the aqueous phase4 or transfer of ClO4− to the organic phase. As has been reported earlier,5 transfer of DMFc+ to the aqueous phase causes the aqueous solution to turn green, which is clearly not the case (Figure 1, flask 1B, upper phase). Thus, oxidation of DMFc is

⊖ − 2ΔTFT W G H+)/2F

(8)

W [E⊖ O2/H2O2]SHE

where is the standard redox potential of the O2/ H2O2 couple in water reported versus standard hydrogen ⊖ electrode, ΔTFT w GH2O2 is the standard Gibbs energy of H2O2 ⊖ transfer from water to TFT, and ΔTFT w GH+ is the standard Gibbs energy of H+ transfer from water to TFT. By taking −1 15 W TFT ⊖ [E⊖ and O2/H2O2]SHE = 0.695 V, Δw GH+ = 50.5 kJ mol , −1 TFT ⊖ Δw GH2O2 = 9.5 kJ mol (see the Supporting Information), TFT ⊖ TFT one calculates [E⊖ O2/H2O2]SHE = 1.169 V. Clearly, [EO2/H2O2]SHE > ⊖ W [E O 2 /H 2 O 2 ] SHE , which means that oxygen reduction is thermodynamically more favorable in the organic phase than in the aqueous phase. This finding supports the homogeneous electron transfer mechanism as being the dominant reaction path. 20014

DOI: 10.1021/acs.jpcc.5b06620 J. Phys. Chem. C 2015, 119, 20011−20015

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(5) Adamiak, W.; Jedraszko, J.; Krysiak, O.; Nogala, W.; HidalgoAcosta, J. C.; Girault, H.; Opallo, M. Hydrogen and Hydrogen Peroxide Formation in Trifluorotoluene−Water Biphasic Systems. J. Phys. Chem. C 2014, 118, 23154−23161. (6) Knochel, P. Modern Solvents in Organic Synthesis; Springer-Verlag: Berlin, Germany, 1999. (7) Thomas, J. C.; Trend, J. E.; Rakow, N. A.; Wendland, M. S.; Poirier, R. J.; Paolucci, D. M. Optical Sensor for Diverse Organic Vapors at ppm Concentration Ranges. Sensors 2011, 11, 3267−3280. (8) Jedraszko, J.; Nogala, W.; Adamiak, W.; Dongmo, S.; Wittstock, G.; Girault, H. H.; Opallo, M. Catalysis at the Room Temperature Ionic Liquid|Water Interface: H2O2 Generation. Chem. Commun. 2015, 51, 6851−6853. (9) Kakiuchi, T. In Liquid-Liquid Interfaces: Theory and Methods; Volkov, A. G., Deamer, D. W., Eds.; CRC Press: Boca Raton, FL, 1996. (10) House, H.; Feng, E.; Peet, N. Comparison of Various Tetraalkylammonium Salts as Supporting Electrolytes in Organic Electrochemical Reactions. J. Org. Chem. 1971, 36, 2371−2375. (11) Shao, Y.; Mirkin, M. V. Probing Ion Transfer at the Liquid/ Liquid Interface by Scanning Electrochemical Microscopy (SECM). Anal. Chem. 1998, 70, 3155−3161. (12) Hu, C.-C.; Liu, K.-Y. Voltammetric Investigation of Platinum Oxides. I. Effects of Ageing on Their Formation/Reduction Behavior as well as Catalytic Activities for Methanol Oxidation. Electrochim. Acta 1999, 44, 2727−2738. (13) Evans, S. A. G.; Elliott, J. M.; Andrews, L. M.; Bartlett, P. N.; Doyle, P. J.; Denuault, G. Detection of Hydrogen Peroxide at Mesoporous Platinum Microelectrodes. Anal. Chem. 2002, 74, 1322− 1326. (14) Huang, J.-F.; Sun, I-W. Formation of Nanoporous Platinum by Selective Anodic Dissolution of PtZn Surface Alloy in a Lewis Acidic Zinc Chloride-1-Ethyl-3-methylimidazolium Chloride Ionic Liquid. Chem. Mater. 2004, 16, 1829−1831. (15) Liu, B.; Bard, A. J. Scanning Electrochemical Microscopy. 45. Study of the Kinetics of Oxygen Reduction on Platinum with Potential Programming of the Tip. J. Phys. Chem. B 2002, 106, 12801−12806. (16) Bard, A. J. In Scanning Electrochemical Microscopy; Bard, A. J., Mirkin, M. V., Eds.; CRC Press: Boca Raton, FL, 2012. (17) Wei, C.; Bard, A. J.; Mirkin, M. V. Scanning Electrochemical Microscopy. 31. Application of SECM to the Study of Charge Transfer Processes at the LiquidLiquid Interface. J. Phys. Chem. 1995, 99, 16033−16042. (18) Burchardt, M.; Wittstock, G. Micropatterned Multienzyme Devices with Adjustable Amounts of Immobilized Enzymes. Langmuir 2013, 29, 15090−15099.

CONCLUSIONS In this work, we have shown that O2 can be reduced to H2O2 in a simple liquid−liquid two-phase system without deliberately adding salts to either the aqueous or the organic phase. Among other two-phase systems with molecular solvents,4,5 the studied system is in fact the simplest one for electron transfer reactions at soft interfaces, as it involves only the reagents and a pair of solvents. To detect H2O2, we performed a series of fluorescence microscopy experiments with a H2O2-sensitive fluorescent dye dissolved in the water phase. Formation of H2O2 was also confirmed by scanning electrochemical microscopy (SECM), where the H2O2 oxidation current was measured in the vicinity of the oil−water interface. Here, to overcome the problem with the electrode surface deactivation during measurements, we applied an electroanalytical procedure where the SECM tip was continuously reactivated. The obtained SECM approach curve was simulated by a finite element model to provide physical insight into the reaction mechanism. The simulations supported the view that the oxygen reduction is a homogeneous reaction in the organic phase: Thermodynamic calculations confirmed that the intrinsic property of the studied system provides favorable conditions for this mechanism. We also measured that ca. 3% of the hydrogen peroxide partitions to the organic phase, and we calculated the transfer energy from water to TFT to be 9.5 kJ mol−1.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpcc.5b06620. Typical cyclic voltammogram and current transient recorded during SECM measurements, details of computer simulations, and determination of H2O2 partition coefficient (PDF)

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AUTHOR INFORMATION

Corresponding Author

*Tel: +48 22 3433306. Fax: +48 22 3433333. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Project “PSPB-035/2010: Electrocatalysis at Droplets” was supported by a grant from Switzerland through the Swiss Contribution to the enlarged European Union.

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REFERENCES

(1) Hess, W. T. In Kirk-Othmer Encyclopedia of Chemical Technology, 4th ed.; John Wiley & Sons: Hoboken, NJ, 2001; Vol. 13, pp 961−995. (2) Goor, G.; Glennebergo, J.; Jacobi, S. In Ullman’s Encyclopedia of Industrial Chemistry; Wiley-VCH: Weinheim, 2012; Vol. 18, pp 393− 427. (3) Campos-Martin, J. M.; Blanco-Brieva, G.; Fierro, J. L. G. Hydrogen Peroxide Synthesis: An Outlook beyond the Anthraquinone Process. Angew. Chem., Int. Ed. 2006, 45, 6962−6984. (4) Jedraszko, J.; Nogala, W.; Adamiak, W.; Rozniecka, E.; LubarskaRadziejewska, I.; Girault, H. H.; Opallo, M. Hydrogen Peroxide Generation at Liquid-Liquid Interface under Conditions Unfavorable for Proton Transfer from Aqueous to Organic Phase. J. Phys. Chem. C 2013, 117, 20681−20688. 20015

DOI: 10.1021/acs.jpcc.5b06620 J. Phys. Chem. C 2015, 119, 20011−20015