Ion effects in water oxidation to hydrogen peroxide

Abstract

We investigate the effect of pH, potential and electrolyte ions in the electrochemical oxidation of water to produce hydrogen peroxide. This process has the potential to provide a low energy-cost route to the generation of hydrogen peroxide, either for in situ use as a “green” oxidant or as part of a water splitting process for the generation of oxygen and hydrogen. Electrodeposited manganese oxide films were used as the working electrode along with aqueous solutions of various ammonium-based cations as the electrolyte. Oxidation of water was carried out at potentials as low as 0.6 V vs. Ag/AgCl at pH 10. Hydrogen peroxide production was found to be highly sensitive to pH, only occurring above pH 9.5, in solutions where the pH had been adjusted by the addition of excess amine. Highly efficient, approaching 100% Faradaic efficiency, production of hydrogen peroxide was observed the range pH 10–10.5. Faradaic efficiency of hydrogen peroxide formation decreased at applied potentials higher than 1 V vs. Ag/AgCl where direct, or further, oxidation to oxygen begins to dominate. Investigation of a number of alkylammonium cations and alkyl sulphate anions of different alkyl chain length indicated that the optimum system is butylammonium sulphate at concentrations around 1 M.

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Introduction

Water oxidation is a vital process, in natural systems as part of both carbon and nitrogen fixation in plants, and also in a variety of important chemical processes. Water oxidation has the potential to offer an energy efficient route to the important and well known oxidant, hydrogen peroxide, which has strong green credentials among oxidants and would be used more widely and efficiently if it could be made more readily on-site or in situ. Equally, water oxidation to oxygen is a key part of efforts to produce hydrogen in large scale from water splitting as an energy storage medium or fuel. Hydrogen is a promising in this regard, as it has a relatively high gravimetric energy density and it can be produced and utilised cleanly by water electrolysis and hydrogen fuel cells respectively.¹ An outstanding problem with the electrolysis of water, however, is the high energy cost of producing hydrogen gas, relative to the energy released by using it as fuel,² i.e., the round-trip energy efficiency is <50%. This is, in part, caused by the high overpotential required in the four-electron, four proton electrochemical oxidation of water to oxygen, (eqn (1)), which is also known as the Oxygen Evolution Reaction (OER).

2H₂O = O₂ + 4H⁺ + 4e⁻, E⁰ = 1.23 V vs. NHE (1)

The same issue arises with concepts involving electrochemical reduction of carbon dioxide to fuels, as the only practical oxidation reaction to couple to this process in large scale is water oxidation.

The utilisation of electrocatalysts to lower this water oxidation overpotential has been widely researched in recent years and has focused on the development and understanding of a number of different catalytic materials. Some of the best known electrocatalysts are metal oxides, the most common and highest efficiency being ruthenium^3,4 and iridium oxides.⁵ Such catalysts that rely on rare and expensive elements are unsuitable for large scale water oxidation, so catalysts that make use of cheaper and more abundant materials⁶ such as cobalt,^7–11 manganese,^12–15 titanium,¹⁶ iron¹⁷ or nickel¹⁸ oxides have been under intense investigation. The influence that chemical composition, crystal structure and nanostructure has on the efficiency of these metal oxide catalysts has also been heavily scrutinised,¹⁶ along with the effects of mixing the metal ions in the oxide.^19–23 However, to date none of these candidates have been able to match the catalytic performance of the precious metal oxides at fundamental turnover frequency and efficiency level.

O₂ is not the only possible product of water oxidation given the complex intermediate steps involved.²⁶ Hydrogen peroxide formation via a two electron, two proton process is an alternative outcome (eqn (2)).

2H₂O → HOOH + 2H⁺ + 2e⁻, E⁰ = 1.776 V vs. NHE (2)

In fact, detailed studies of the water oxidation centre in Photo-system II, which involves a manganese–oxo cluster, have been instructive in this respect. These have deepened the understanding of the mechanisms of water oxidation, providing some insight into the details of the four, distinct oxidation and proton transfer steps.²⁴ Theoretical modelling shows that the formation of the first O–O bond, i.e., to form a per-oxy species via the loss of the first two electrons and two protons, is a key step in the overall process²⁵ and that the loss of the third electron may present the greatest energy barrier.

Similarly, Ando et al. have also suggested that a 2-electron reaction of water oxidation to H₂O₂,²⁷ may be possible in vitro. Since E⁰ for hydrogen peroxide production is substantially higher than that for oxygen, the ideal catalyst/electrolyte combination for this process will be one that lowers the overpotential, and possibly also the E⁰(H₂O/H₂O₂), of eqn (2), thereby promoting the preferential formation of hydrogen peroxide over oxygen. However, in practical water oxidation research, less emphasis has been placed on the interaction between the catalysts and the electrolyte surrounding the catalytic site. Clearly proton activity and proton activity (pH) buffering are key aspects of the role of the electrolyte, both from thermodynamic and kinetic points of view. There is also a potential for specific chemical interaction aspects of the electrolyte to influence the mechanism of the reaction. In this direction, evidence was found by Izgorodin et al.²⁸ that suggests that direct oxidation of water to hydrogen peroxide preferentially occurs when alkylammonium cation based electrolyte solutions are used in conjunction with a manganese oxide electro-catalyst.²⁸ It was hypothesised that the process involved preferential solvation of the hydrogen peroxide product by the alkylammonium cation and/or its conjugate base, lowering the free energy of the solvated species and thereby lowering the potential at which it is formed. Hydrogen peroxide efficiencies as high as 77% at 0.59 V vs. Ag/AgCl were demonstrated in an electrolyte that consisted of 1 M of a butylammonium sulfate salt/base mixture at pH 10. These are relatively low potentials for a water oxidation process compared to the potentials typically required for direct water oxidation to oxygen at the same pH. It was suggested that this process could be used as a means of direct production of hydrogen peroxide for use as an oxidant. Alternatively, if combined with a subsequent decomposition to form oxygen (eqn (3)), the process appears to offer a potentially low energy-cost alternative to the traditional 4-electron oxidation oxygen evolution reaction.

2HOOH → H₂O + O₂ (3)

Hydrogen peroxide is used as a strong oxidant and bleaching agent and is most commonly produced using the anthroquinone auto-oxidation process, which is energy intensive and can have environmentally unfriendly byproducts.²⁹ Small scale H₂O₂ production is an active area of research^30–39 as it could eliminate the need for expensive handling and transportation of H₂O₂ from the relatively small number of production facilities world-wide. It also opens up possibilities for electro-generated H₂O₂ to be generated in situ for important chemical processes⁴⁰ or even as an alternative standalone fuel in a hydrogen peroxide fuel cell.^41–44

Based on the original hypothesis of Izgorodin et al.²⁸ that solvated hydrogen peroxide complexes were responsible for their observation of preferential formation of H₂O₂ we investigate in the present work the role of the electrolyte cations and anions involved. The organic salts examined include a number that originate from the protic ionic liquid family of salts, being those that in their pure state have melting points below 100 °C; in the present work this ionic liquid nature serves to increase solubility in the 1 M aqueous solutions that are used. We also investigate the proton activity dependence of the amount and efficiency of H₂O₂ produced in these electrolytes, with the ultimate goal of further optimising the features of this process.

Experimental

Catalytic film deposition

Manganese oxide catalyst films were electrodeposited using the method described by Zhou et al.⁴⁵ Films were electrodeposited on gold in a 2 mL solution of 1 M ethylammonium nitrate and 10 mM manganese acetate. The electrodeposition was conducted at room temperature using chronopotentiometry at a constant current density of 200 μA cm⁻² over 10 minutes. The films were then heat-treated at 90 °C for 30 minutes by placing the electrode on a hotplate. Manganese acetate tetrahydrate 99.99% was purchased from Sigma-Aldrich. 70% nitric acid was purchased from Univar.

Electrolyte preparation

Electrolyte solutions were made by dissolving a sufficient amount of the relevant acid in each case to achieve 1 M acid in water. The required amine was added to adjust the pH using a TPS smartCHEM-pH meter to the desired value in the range between 9 and 11. This is considerably beyond the simple 1 : 1 salt mole ratios in each case and in fact is approximately the buffer region of the base in each case. This approach to the electrolyte preparation, using the glass electrode based pH measurement ensures that constant proton activity is obtained across all of the systems.

The amines used were 99.5% butylamine and diethylamine and 99% hexylamine purchased from Sigma-Aldrich. The acids used were 98.08% sulphuric acid purchased from Unilever, 99.5% methanesulfonic acid and 99% ethanesulfonic acid and toluenesulfonic acid monohydrate purchased from Sigma-Aldrich.

Electrochemistry

All electrochemical measurements were carried out using a multi-channel potentiostat (VMP2, Princeton Applied Research). Electrochemical cells consisted of the MnO_x film as the working electrode, a Ag/AgCl reference electrode and a titanium mesh counter electrode. When measuring cyclic voltammograms of the different electrolytes with various ions and pHs, 0.5 mL of the electrolyte and a scan rate of 1 mV s⁻¹ was used. Conductivity was determined by potentio-electrochemical impedance spectroscopy (PEIS). The solution resistance of the electrolytes between two flat gold electrodes 1 cm apart was used to calculate the conductivity.

The measurements of H₂O₂ production were carried out using potentiostatic coulometry where a constant potential was applied to the cell and the current was measured over a known time. In these measurements a fritted compartment was used to contain the counter-electrode in order that back reduction of the H₂O₂ could not take place at the cathode. The total charge that passed through the cell was then calculated using the VMP software by integrating the area under the current vs. time curve. The electrolyte was then transferred to a separate vessel for chemical determination of the H₂O₂ content. The number of moles of H₂O₂ in solution was determined using a standard titration of the electrolyte using KMnO₄ according to the reaction:

2MnO⁴⁻ + 5H₂O₂ + 6H⁺ → 2Mn²⁺ + 5O₂ + 8H₂O (4)

For this titration 250 μL of 1 M sulphuric acid was added to 250 μL of the electrolyte. A 7.55 × 10⁻³ M potassium permanganate solution was used as the titrant and added to the analyte in volumes of 10 μL. Additions were separated by 20 minutes to allow the absorption reading to stabilise. A UV-Vis spectrometer measured absorption at 565 nm wavelength. The Faradaic efficiency of H₂O₂ production was then determined using the total charge passed through the cell during the experiment. An efficiency <100% in this respect reflected the competing production of oxygen in the oxidation reaction. Errors in calculations were estimated from repetitions (n = 4–5) in selected cases using the same film. A single film was used to carry out each series of experiments (i.e. a different film for each electrolyte, one film used for all of the comparisons of pH and one film used for time measurements). This was done to maintain consistency within the experiments and allow for any degradation of the films with use. The potassium permanganate was purchased from Sigma-Aldrich and the concentration of the KMnO₄ solution was determined by titration with sodium oxalate. The total amount of H₂O₂ produced was divided by the total amount of Mn in the film to determine the turn over number of an experiment and this number was then divided by the length time of applied potential to determine the turn over frequency. The total amount of Mn in the film was estimated by calculating the charge used in the deposition and equivocating each mole of Mn deposited to 2 moles of electrons. This represents an upper bound on the amount of Mn in the film as some charge is lost during deposition to other processes.

Results

The cyclic voltammograms in Fig. 1 show an oxidation peak beginning at around 0.6 V vs. Ag/AgCl at pH > 10, before the main water oxidation to oxygen peak appears at 1.3 V vs. Ag/AgCl. A clear correlation is revealed in Fig. 2 between this oxidation peak at 0.6 V vs. Ag/AgCl and the start of H₂O₂ production in the butylammonium sulfate (BAS), electrolytes. This is consistent with the previously reported oxidation potential using this type of electrolyte.²⁸ This indicates that the water oxidation to hydrogen peroxide reaction starts to occur at potentials considerably lower than that for oxygen production and that the hydrogen peroxide produced is sufficiently stable in the solution for subsequent determination by the chemical means used. Fig. 1 also shows the influence that the pH of the electrolyte has on the performance of the water oxidation processes. From the cyclic voltammograms in Fig. 1, the first oxidation peaks start to appear when using electrolytes with a pH higher than ∼9.5. When the pH approaches 11 this peak disappears while the second peak (i.e. the OER process) shows the highest current of all of these electrolytes, indicating the strong influence that pH has on the OER reaction on this electrocatalyst.

Fig. 1 Cyclic voltammograms during water oxidation in 1 M BAS electrolyte, at various pHs on a MnO_x electrode. The scan rate was 1 mV s⁻¹.

Fig. 2 H₂O₂ production at 0.8 V vs. Ag/AgCl during controlled potential electrolysis for 10 minutes on an MnO_x electrocatalyst in 1 M BAS electrolyte at various pH values. (0.8 V vs. Ag/AgCl over 10 min).

In the controlled potential electrolysis runs in Fig. 2, detectable H₂O₂ production begins around pH 9 and tends to increase as the pH is increased. (Note that Fig. 1 shows that the current is relatively constant around 0.8 V vs. Ag/AgCl in all cases; therefore the shift in E⁰ due to pH changes in Fig. 2 is not a significant factor in determining the current). The Faradaic efficiency with respect to H₂O₂ production reaches 100% in the pH region between 10 and 10.4. At higher pH, Faradaic efficiency decreases as a result of further oxidation on the electrode to generate oxygen. Experiments were not conducted at even higher pH values because the addition of more amine caused the electrolyte to become too volatile to maintain a constant pH. At the lower pHs, when the electrolyte is closer to neutral, the MnO_x films are known to lose their effectiveness as catalysts.^15,28 Thus, it appears that this particular process of water oxidation to H₂O₂ occurs most readily within a specific pH range around 10–11. Referring to the pK_a data in Table 1, this corresponds approximately to the buffer region in each case, suggesting that a buffering action in absorbing the protons produced in the reaction may be a factor in the efficiency of this process. However, as was shown by Izgorodin et al.,²⁸ inorganic buffer solutions do not support H₂O₂ production and therefore buffering action and/or pH alone are not the origins of this mechanism.

Table 1 H₂O₂ production data for 1 M ammonium-based electrolytes at pH 10 for

Electrolyte	Conductivity (S cm^−1)	Current density at 0.8 V vs. Ag/AgCl (mA cm^−2)	Moles H2O2 produced (10 min, at 0.8 V vs. Ag/AgCl) (μmol) (±0.04)	% Faradaic efficiency of H2O2 production (10 min, at 0.8 V vs. Ag/AgCl) (±6%)	pKa (base)
BAS	0.11	0.2	0.37	106	10.77
BAMeS	0.06	0.13	0.11	62	10.77
BAEthS	0.04	0.12	0.27	101	10.77
DEAS	0.04	0.07	0.11	60	11.2

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Fig. 3(a)–(d) show cyclic voltammograms and the associated trends in H₂O₂ production and efficiency from the potentiostatic coulometry experiments, for a range of electrolytes, including butylammonium methanesulfonate (BAMeS) and butylammonium ethanesulfonate (BAEthS), as the potential applied to the cell is increased. The pH of the solution is maintained at 10 in each case and the concentration of the electrolyte is 1 M (with respect to the anion). Cyclic voltammograms of 1 M butylammonium tosylate, diethylammonium tosylate and hexylammonium sulfate electrolyte solutions (pH 10) were also measured, but were found to have relatively low currents at the expected potentials for water oxidation (see ESI Fig. A†). A hexylammonium tosylate electrolyte was also tested but upon addition of the amine to acid a solid formed that was insoluble in water.

Fig. 3 (a–d): H₂O₂ production during 10 minutes constant potential electrolysis in (a) BAS, (b) BAMeS, (c) BAEthS and (d) DEAS electrolytes at 1 M pH 10 in each case. The scan rate for the cyclic voltammetry was 1 mV s⁻¹.

In the BAS and BAMeS systems the amount of H₂O₂ produced seems to plateau around 1 V vs. Ag/AgCl before increasing again at higher potentials when the second oxidation process occurs. This tends to follow the current density in the cyclic voltammograms such that a higher current in this region means more H₂O₂ generated. In the BAEthS system there is a more significant drop in the amount of H₂O₂ as the applied potential is increased from 0.8 V to 1.0 V vs. Ag/AgCl. The cell using the 1 M BAS electrolyte shows twice the amount of H₂O₂ produced at all potentials, as compared to the BAMeS and DEAS cells.

Fig. 3d shows the production and efficiency of the cell when the secondary amine based electrolyte, DEAS, is used. In this case the cyclic voltammogram shows two oxidation peaks that occur at higher potentials than the other cells, at 0.7 and 1.0 V vs. Ag/AgCl respectively, before the main water oxidation to oxygen (OER) peak at 1.3 V vs. Ag/AgCl. This additional process may be the appearance of the amine oxidation process to the N-oxide which occurs more readily in secondary and tertiary amines than in primary amines. It has been demonstrated previously that the butylammonium sulphate systems are not oxidised under these conditions.²⁸

Also shown in Fig. 3 in each case is the apparent Faradaic efficiency calculated as described in the experimental section (the uncertainty in the efficiency measurement at the lower applied potentials, 0.7 V and below, makes further analysis difficult in these cases). In the BAS electrolyte, the H₂O₂ production efficiency reaches 100% within error around 1.0 V vs. Ag/AgCl, while in the other electrolytes the efficiency is lower. The BAS result achieved here is somewhat higher than previously reported by Izgorodin et al.,²⁸ due to the different potential range studied in that work; it appears from the present work that potentials around 0.9–1.0 V vs. Ag/AgCl are optimum for efficient H₂O₂ production.

Generally the efficiency of hydrogen peroxide production drops as the potential rises above 1.2 V vs. Ag/AgCl. This is as expected due to the consumption of current by the competing OER that begins to occur at these higher potentials. It is also possible that the further 2 electron oxidation of the H₂O₂ begins to occur at these potentials.

Table 1 compares the data from Fig. 3 at 0.8 V vs. Ag/AgCl showing the current density, the quantity of H₂O₂ produced and the Faradaic efficiency at this potential for each electrolyte. It can be seen that BAS and BAEthS have the highest Faradaic efficiency. It also appears, comparing the butylammonium and diethylammonium cations, that the former produces a significantly larger amount of H₂O₂. The variation in efficiency across the salts in Fig. 3 indicates an involvement of the cation and/or anion in promoting the high efficiency H₂O₂ formation route. To explore this further we list in Table 1 the conductivities of the electrolytes; it appears that, although there are significant differences between them, the conductivity of the electrolyte does not seem to have a strong influence over the amount of H₂O₂ produced. Although the BAS electrolyte has the highest conductivity and also produces the most H₂O₂ at 0.8 V vs. Ag/AgCl, the BAEthS electrolyte has a low conductivity, but produces the second highest amount of H₂O₂ at the same potential. Higher conductivity may result in higher current densities at a particular potential (as shown in the third column of Table 1). However, it is likely that other factors such as the rate of H₂O₂ disproportionation in the different electrolytes have a greater influence over the detected H₂O₂.

Fig. 4 shows the change in the amount of H₂O₂ produced as the water oxidation is conducted over different lengths of time. This experiment has been carried out both with and without a separator isolating the working and counter electrodes in the cell. In the two compartment cell, an almost linear increase in the amount of H₂O₂ produced is observed over two hours.

Fig. 4 Production of H₂O₂ in 1 M pH 10 BAS electrolyte when potential is applied to the cell for various lengths of time (cell volume 500 μL of 1 M electrolyte).

A clear limit of about 1 μmol is reached after one hour of applied potential in a cell without a separator, indicating reduction on the counter electrode of the H₂O₂ produced and that therefore compartment separation is important in respect of the efficiency of this process at longer electrolysis times. However, flow through type cells that might be used in an in situ application of this process may not require separation because of the relatively shorter residence time. While it is likely that only manganese species located on the surface of the film are catalytically active it is nonetheless instructive to calculate the turnover number (TON) and turnover frequency (TOF) on the basis of total manganese in the film. On this basis, the maximum TON observed thus far in our experiments without any serious sign of degradation is 20 over 2 hours. It is important to note that this represents a lower bound on the real number since some of the manganese is likely to be inactive. The TOF calculated on the same basis can be as high as 0.378 mol (H₂O₂) mol (Mn)⁻¹ min⁻¹ (from experiments where pH = 10).

Discussion

The results shown above indicate that the pathway to hydrogen peroxide lowers the potential of water oxidation in this electrolyte. Energy calculations performed by Rossmeisl et al.⁴⁶ for a rutile-type oxide catalyst examined the following water oxidation reaction mechanism:

The calculations suggest that the third step, which produces the HOO˙ species, is the least thermochemically favourable step (i.e. highest potential required) and is thus thought to be the potential-limiting reaction. This is believed to not depend upon the type of catalytic surface. Dau et al.²⁶ described a related series of reactions that can instead lead to the formation of hydrogen peroxide in the second step in an alkaline electrolyte (where ∑ is a catalytically active surface site):

∑ + OH⁻ → ∑–OH + e⁻ (9)

∑–OH + OH⁻ → ∑ + H₂O₂ + e⁻ (10)

As was demonstrated in energetic calculations by Izgorodin et al.²⁸ the H₂O₂ formed in this reaction (eqn (10)) is likely to be hydrogen bonded to the free amine in the electrolyte. This may aid the breaking of the ∑–OH bond in this reaction scheme, leading to the net formation of solvated H₂O₂ at low potentials without further oxidation to O₂. This is also consistent with the observation of a lower H₂O₂ production rate seen when the 1 M DEAS electrolyte is used, since a secondary amine would offer more limited hydrogen bonding.

Examining the pH effect on these reactions, Takashima et al.¹⁵ hypothesised that the presence of Mn³⁺ in MnO_x catalytic films is a precursor for water oxidation. Their work focused on the dramatic effect that pH has on electrochemical water oxidation on MnO₂ films, and showed that the reaction was unfavourable at pHs at which Mn³⁺ is unstable. For pH values >9, the com-proportionation of Mn²⁺ and Mn⁴⁺ in the film stabilises the Mn³⁺ centre on which water oxidation occurs. Thus there may be multiple origins of the pH effects seen in the present work; nonetheless the presence of the organic cation/amine appears to be strongly implicated in the mechanism involved.

Summary and conclusions

The interactions of various ammonium-based, free amine containing salt solutions with electrocatalytic MnO_x films in a water oxidation cell have been examined. A correlation between the oxidation peaks in the cyclic voltammograms of these cells and the production of H₂O₂ was shown, beginning at 0.6 V vs. Ag/AgCl which yields a relatively low potential of 0.2 V compared to the equilibrium potential of a standard OER. The production of H₂O₂ was found to be sensitive to the cation and anion of the salt as well as the pH of the electrolyte. The amount of H₂O₂ produced was found to increase linearly with time when the counter electrode is isolated from the rest of the cell over a period of several hours. This creates clear potential for the H₂O₂ to be used as an in situ generated oxidant. Further work in this demonstrating such applications are underway and will be reported in due course.

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Footnote

† Electronic supplementary information (ESI) available: Cyclic voltammograms of low performance electrolytes and description of catalytic MnO_x films. See DOI: 10.1039/c4ra05296j

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