Temperature-Dependent Kinetic Studies of the Chlorine Evolution

Feb 17, 2017 - Iman Sohrabnejad-Eskan,. † ... of Inorganic Materials Chemistry, Department of Chemical Engineering and Chemistry, Eindhoven Universi...
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Temperature-Dependent Kinetic Studies of the Chlorine Evolution Reaction over RuO(110) Model Electrodes 2

Iman Sohrabnejad-Eskan, Andrey E. Goryachev, Kai Steffen Exner, Ludwig A. Kibler, Emiel J. M. Hensen, Jan Philipp Hofmann, and Herbert Over ACS Catal., Just Accepted Manuscript • DOI: 10.1021/acscatal.6b03415 • Publication Date (Web): 17 Feb 2017 Downloaded from http://pubs.acs.org on February 20, 2017

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Temperature-Dependent Kinetic Studies of the Chlorine Evolution Reaction over RuO2(110) Model Electrodes Iman Sohrabnejad-Eskan1, Andrey Goryachev2, Kai S. Exner1,3, Ludwig A. Kibler3, Emiel J. M. Hensen2, Jan P. Hofmann2, Herbert Over1* 1

Physikalisch-Chemisches Institut, Justus-Liebig-Universität, Heinrich-Buff-Ring 17, 35392 Gießen, Germany 2 Laboratory of Inorganic Materials Chemistry, Department of Chemical Engineering and Chemistry, Eindhoven University of Technology, P.O. Box 513, 5600MB Eindhoven, The Netherlands 3 Institut für Elektrochemie, Universität Ulm, Albert-Einstein-Allee 47, 89081 Ulm, Germany * Corresponding author: E-mail: [email protected]

Abstract Ultrathin single crystalline RuO2(110) films supported on Ru(0001) are employed as model electrodes to extract kinetic information about the industrially important chlorine evolution reaction (CER) under well-defined electrochemical conditions and variable temperatures. A combination of chronoamperometry (CA) and on-line electrochemical mass spectrometry (OLEMS) experiments provides insight into the selectivity issue: At pH = 0.9 the CER dominates over oxygen evolution, whereas at pH = 3.5 oxygen evolution and other parasitic side reactions contribute mostly to the total current density. From temperature-dependent CA data for pH = 0.9 we determine the apparent free activation energy of the CER over RuO2(110) to be 0.91 eV, which compares well with the theoretical value of 0.79 eV derived from first principles microkinetics. The experimentally determined apparent free activation energy of 0.91 eV is considered as a benchmark for assessing future improved theoretical modeling from first principles.

Keywords: chlor-alkali electrolysis, chlorine evolution reaction (CER), oxygen evolution reaction (OER), RuO2, apparent free activation energy, selectivity

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1. Introduction Chlorine manufacture via chlor-alkali electrolysis is the second largest industrial electrochemical process with an annual production of more than 70 million tons Cl2 requiring a total electric energy of more than 200 TWh.1-3 The membrane method is current state of the art, where a Nafion type ion exchange membrane is employed to separate the anodic and cathodic half cells. Chlor-alkali electrolysis is carried out at 80°C in a 5 M concentrated NaCl (brine) solution containing hydrochloric acid producing Cl2 in the anodic compartment and NaOH as well as H2 in the cathodic compartment.4 Advanced electrolyzers return the produced H2 to a kind of integrated fuel cell (called gas diffusion electrodes) to form water, thereby reducing the electric power consumption by 30%.5 In the potential region of chlorine evolution reaction (CER, U0CER = +1.36 V vs. SHE under standard conditions), the formation of gaseous oxygen at the anode (OER: oxygen evolution reaction, U0OER = +1.23 V vs. SHE under standard conditions)4 is thermodynamically preferred. This results in Faradaic efficiency losses as well as the need for costly product purification steps. At higher pH, hypochlorite and hypochlorous acid are able to be formed by Cl2 disproportionation or by the direct oxidation of chloride anions at the anode when for instance OH− leaks through the membrane into the anodic compartment.6,7 For the CER, the reversible electrode potential is pH independent (but dependent on the Cl− and Cl2 concentration), in contrast to the OER where the reversible electrode potential increases with decreasing pH. Therefore, the selectivity can be shifted towards CER by keeping the pH of the electrolyte solution low in order to raise UOER, while adjusting the concentration of Cl− as high as possible so that UCER is low. In the membrane chlor-alkali process, the lowest possible pH is about 2, since the Nafion membrane is unstable in solutions with higher proton activities. The industrially used CER catalyst in form of Dimensionally Stable Anodes (DSA®) consists of Ti plates coated with rutile TiO2-RuO2 mixed oxides (typically in a 70:30 molar ratio).8-12 For this reason RuO2(110), as the most stable surface termination of rutile RuO2, is envisioned as an appropriate single-crystalline model electrode for investigating the activity and selectivity issue of CER. The corresponding single-crystalline RuO2(110) model anode can be grown epitaxially on single crystal Ru(0001) substrates under ultra-high vacuum (UHV) conditions.13 Recent theoretical studies resolved the reaction mechanism and the complete free energy profile along the reaction coordinate for the CER at the RuO2(110) model electrode.14 The same authors suggested that the apparent free activation energy for the CER on RuO2(110) could be derived -2ACS Paragon Plus Environment

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from temperature-dependent experiments and compared to the theoretical free energy diagram.15 This may offer a powerful tool to validate current first principles studies in electrocatalysis. In heterogeneous catalysis of gas phase reactions, temperature-dependent experiments in the form of Arrhenius plots are the standard approach to determine apparent activation energies, whereas this kind of experiment is challenging in electrocatalysis and therefore has rarely been carried out.16,17 Temperature variations affect the viscosity of the electrolyte solution (transport phenomena), the reversible half-cell potential UCER (thermodynamics), the exchange current density j0 (kinetics), the potential of the reference electrode and the selectivity of the competing OER and CER so that the measured current density as a function of the electrode potential depends on the temperature in a complex way. However, the apparent free activation energy enters only the exchange current density j0 that needs to be extracted from the measurements. In the present study, we address the temperature-dependent activity and selectivity of a RuO2(110) model electrode in the CER under ultra-clean electrochemical conditions by employing on-line electrochemical mass spectrometry (OLEMS)18 and chronoamperometry (CA). OLEMS allows for a fast detection of dissolved gaseous products (e.g. Cl2) close to the model electrode surface. The apparent free activation energy for the CER over RuO2(110) is obtained from temperature-dependent experiments and amounts to 0.91 eV which compares well with the apparent free activation energy derived from the free energy profile along the reaction coordinate as calculated by first principles kinetics (0.79 eV).

2. Experimental Section 2.1 RuO2(110) preparation and characterization The ultrathin RuO2(110) film was grown epitaxially on a 8 mm disk-shaped Ru(0001) single crystal (MaTecK, Jülich, Germany) under ultrahigh vacuum (UHV) conditions. First, the Ru(0001) single crystal was sputtered with Ar+ ions at the temperature of 380°C for 20 minutes in order to clean the sample. Subsequently, the sample was heated up to 780°C to smoothen the rough surface. This procedure was repeated several times until the low energy electron diffraction pattern (LEED) showed an intensive hexagonal diffraction pattern (with low background) corresponding to a clean Ru(0001) surface. Next, the sample was annealed at 780°C for 20 min in an oxygen gas atmosphere of p(O2) = 2·10−7 mbar to deplete the surface near region from carbon contamination. -3ACS Paragon Plus Environment

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Figure 1: Long-term oxidation of Ru(0001) by dosing 3·10−5 mbar of O2 at a sample temperature of 380°C. After 60 minutes, besides oxide reflections of RuO2(110) substrate reflections of Ru(0001) are still strongly visible (blue crosses). After 120 minutes exposure, the substrate reflections have vanished, indicating that the Ru(0001) surface is fully covered with RuO2(110). RuO2(110) on Ru(0001) was prepared at a temperature of 380°C with an oxygen pressure of 3·10−5 mbar for 120 min.13 In Figure 1, a series of LEED patterns shows the evolution of the RuO2(110) oxide film as a function of exposure time until the Ru(0001) surface is fully covered by RuO2(110).

2.2 Electrochemical cell and On-Line Electrochemical Mass-Spectrometry (OLEMS) For electrochemical experiments, a standard water jacket electrochemical glass cell was used in a three-electrode configuration (Figure 2a). The cell was equipped with a thermostat (MGW Lauda) and two argon gas lines for purging the cell through and over the electrolyte. Before each experiment, the electrochemical cell was cleaned with chromic acid (20 g K2Cr2O7 in 400 mL H2SO4 conc.) and rinsed with copious amounts of Milli-Q water (18.2 MΩ cm). As working electrode, a RuO2(110)/Ru(0001) single crystal (prepared by the procedure described above) was mounted into the PEEK sample holder. A platinum plate (2 x 2.5 cm2) and a Red Rod electrode (Radiometer Analytical, saturated KCl, E0RE = +0.215 V) were employed as counter and reference electrodes, respectively, and connected to an Ivium Compactstat potentiostat (Ivium Technologies). For gaseous product analysis, on-line electrochemical mass spectrometry (OLEMS) was used. OLEMS is a variation of differential electrochemical mass spectrometry (DEMS)19-21 that brings a small capillary with a semipermeable Teflon membrane very close to the electrode surface to improve the sensitivity. The design and detailed description of the OLEMS system used can be found in the paper of Wonders et al.18 An OLEMS capillary made -4ACS Paragon Plus Environment

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of PEEK containing a porous Teflon plug was fixed in a quartz tube and approached to 20 µm above the surface of the working electrode utilizing a video microscope. The OLEMS tip was cleaned by rinsing it several times with ethanol and Milli-Q water (18.2 MΩ cm) prior to mounting it to the EC cell. The OLEMS capillary was connected to a quadrupole mass spectrometer (Balzers Quadstar, Prisma QME 200) via a two-step differential pumping system in order to achieve an operating pressure in the ionization chamber of the QMS of approximately 5·10−7 mbar. Signals of m/z = 36 (corresponds to H35Cl+ with extra H stemming from water dissociation inside the ion source of the MS), m/z = 32 (O2+) and the background pressure were constantly monitored and recorded via a 4-channel AD converter interface of an Ivium potentiostat (Ivium Technologies, Peripheral Port Expander). All potentials reported in this work are plotted versus standard hydrogen electrode (SHE). The presence of HCl (H35Cl+, m/z = 36) rather than Cl2 signals in the mass spectrometer points to dissociation of Cl2 and recombination with H from water, present as the main component in the MS chamber. The employed porous Teflon plugs in the OLEMS capillary always leak a certain amount of water into the MS system. Large amounts of HCl and Cl2 in the MS chamber can lead to changes of the sensitivity of the MS for other species, such as O2, hampering their quantification.22

2.3 Temperature-dependent chlorine evolution reaction (CER) For the investigation of the chlorine evolution reaction (CER), a 5 M NaCl solution in 10−2 M HCl was prepared as electrolyte solution. The temperature of the electrolyte solution was kept constant at one of the selected temperatures (T = 25, 30, 35, 40°C) using a thermostat. Prior to the experiments, the electrochemical cell was cleaned as described above and the electrolyte was purged with Ar gas for 1 hour. All reagents were of ultra-pure grade and used without additional purification.

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Figure 2: a) Scheme of the OLEMS experimental setup including the three-electrode electrochemical cell equipped with an OLEMS capillary and a magnified view on the OLEMS tip sample configuration. b) Schemes of the CER pulse experiments conducted in 5 M NaCl + 10−2 M or 5 M NaCl + 10-5 M HCl at temperatures T = 25 – 40°C. Each potential pulse was kept for 10 s and repeated once, while between the pulses the electrode potential was kept at a rest potential of 1 V for 10 minutes. The applied electrode potential is increased in steps of 10 mV. After each experiment, a CV of the RuO2(110) model electrode was taken. Abbreviations: WE: working electrode, CE: counter electrode, RE: reference electrode, PKR: pressure gauge for background pressure in MS chamber.

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As indicated in Figure 2b, for each selected temperature the RuO2(110) electrode was subjected to a sequence of potential pulses starting from the CER pre-polarization regime (U = +1.28 V, 10 s) with an intermediate resting potential (URES = +1.00 V, 10 min) at which no current was observed during the cyclic voltammetry (CV) scans. It is imperative to retain with the same model electrode RuO2(110)/Ru(0001) and the same WE/OLEMS tip geometry in order to be able to compare the OLEMS signals for various temperatures directly. A time period of 10 min at the resting potential in between consecutive CER pulses is necessary to remove produced Cl2 from the near-electrode region and to separate individual OLEMS signals of the evolved products. At the starting potentials where no Cl2 evolution was observed, the duration of the resting potential was set to only 2 min. For reproducibility reasons, each pulse was repeated twice. Subsequently, the applied electrode potential U was increased in 10 mV steps until the electrode potential reached a value of +1.40 V. Prior to the next experiment, the EC cell could equilibrate to the chosen temperature within a time period of 2 hours. The reference electrode potential was corrected with respect to the applied temperature (cf. Electronic Supporting Information (ESI), Figure S1a). CVs were recorded before and after each experimental sequence. Besides the electrochemical data (chronoamperometry, cyclic voltammetry), OLEMS signals of O2 (m/z = 32) and HCl (m/z = 36) were continuously recorded; the background pressure inside the QMS was also monitored. The masses were chosen based on the mass spectra recorded during CER (cf. ESI, Figure S1b) Additionally, a similar sequence of temperature-dependent experiments was conducted at higher pH, using a 10−5 M HCl solution in 5 M NaCl. A sequence of potential pulses in the range from +1.28 V to +1.41 V was applied stepwise in 10 mV steps. The resting potential of +1.00 V was applied in-between the pulses with a duration of 10 min. To compare the temperature dependence at less acidic pH values with those at higher pH, two temperatures, 25°C and 35°C, were selected. In order to compare the experimental investigations with a recent full kinetics study from first principles,14 the equilibrium potential of CER at 25°C was determined based on the measured current density in the polarization regime of CER (U < 1.29 V vs. SHE) as function of the applied electrode potential in Figure S2 (cf. ESI) and linearly extrapolated to zero current density. In this way, the equilibrium potential of CER was determined to be UCER = 1.26 V vs. SHE for 25°C, which is consistent with a previous study.23 Guerrini et al. assumed an activity of -7ACS Paragon Plus Environment

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Cl2 and Cl− of 0.01 and 5.01, respectively, from which they derived an equilibrium potential of CER of 1.26 V vs. SHE applying Nernst’s equation 0 UCER (T , a(Cl − ), a(Cl2 )) = UCER (T ) −

kBT kT ⋅ ln a(Cl − ) + B ⋅ ln a(Cl2 ) e 2e

(1).

In equation (1), U0CER denotes the equilibrium potential of CER under standard conditions (U0CER = 1.36 V vs. SHE), while kB, e, and T denote Boltzmann’s constant, the absolute temperature in Kelvin and the elementary charge of an electron, respectively. The CER equilibrium potentials for higher temperatures were calculated using Nernst’s equation (1). Equation (1) indicates that the equilibrium potential of CER is independent of pH. In the SI, Figure S3, we demonstrate that a variation of the pH value does also not affect the kinetics of the underlying CER over a RuO2(110) electrode as previously reported by Guerrini et al.23 and Consonni et al.27

3. Experimental Results 3.1 Experiments at pH = 0.9 In the first series of measurements, CER over a single-crystalline RuO2(110) model electrode was studied in a 10−2 M HCl (ultra-pure) + 5 M NaCl (ultra-pure) solution. The pH of this electrolyte solution was measured with a glass electrode to be 0.9.24 In the pulsed chronoamperometric measurements the RuO2(110) electrode was subjected to a sequence of potential pulses with increasing electrode potential in steps of 10 mV starting from U = +1.26 V and terminating at U = +1.40 V vs. SHE (cf. Figure 3). As advantage of OLEMS, the produced chlorine can be detected even in very small quantities, long before the first bubbles of Cl2 evolving from the WE surface can be visually observed. While in case of T = 25°C (cf. Figure 3a) the onset of the ion current of m/z = 36 (HCl) is measured at U = +1.38 V vs. SHE, we detect the initial chlorine evolution at U = +1.37 V vs. SHE and at U = +1.36 V vs. SHE for T = 30°C and T = 35°C, respectively (cf. Figure 3b,c). In the corresponding potential range up to U = +1.40 V vs. SHE, the produced oxygen gas is below the detection limit of OLEMS. Therefore, the measured electric current can almost exclusively (by more than 95%) be ascribed to the CER. For the OLEMS experiment at T = 40°C we needed to change the OLEMS tip with the consequence that the mass spectrometry data for T = 40°C is not comparable (different geometry of the tip to the electrode surface) to the preceding OLEMS experiments at lower temperatures and therefore has been omitted. -8ACS Paragon Plus Environment

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Besides the OLEMS data, Figure 3 displays the measured current densities by chronoamperometric (CA) experiments as a function of the electrode potential for various temperatures. The CA measurements are more sensitive than OLEMS (the electrode potentials for onset current in CA is substantially lower than for the onset OLEMS signal), but in contrast to OLEMS not element specific. It turns out that with increasing electrode potential the current density raises first linearly and then exponentially as expected from the Butler-Volmer equation.25

Figure 3: Chronoamperograms (CA) of RuO2(110)/Ru(0001) working electrode in 10−2 M HCl and 5 M NaCl subjected to a series of electrode potential pulses (values given w.r.t. SHE) recorded at a) T = 25°C, pH = 0.9; b) T = 30°C, pH = 0.9; c) T = 35°C, pH = 0.9; d) T = 40°C, pH = 0.9; OLEMS experiments: Ion currents of m/z = 36 (HCl, blue line) recorded at a) T = 25°C; b) T = 30°C; c) T = 35°C.

In order to be able to extract the temperature dependence of the exchange current density j0, we need to adjust the actual reaction conditions in a way that we can neglect the T-dependence of the transport properties and of the OER/CER selectivity. Therefore, we restrict ourselves to a potential region where kinetics is dominating over transport (keeping the electrode potential below 1.40 V). Secondly, we analyze the current density data only for pH = 0.9 so that the CER strongly dominates the electrolysis reaction in the considered potential range. -9ACS Paragon Plus Environment

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In general, with increasing temperature the current density increases for fixed electrode potential (cf. Figure 3). This finding is due mainly to two effects: firstly, the equilibrium potential of CER (UCER) is temperature-dependent in that with increasing temperature UCER shifts to smaller values (Nernst equation (1)). Therefore, the overpotential ηCER is a function of temperature, as the overpotential is defined as the difference between the applied electrode potential pulse in Figure 3 and UCER; we may recall that the overpotential ηCER constitutes the driving force for the CER. In order to analyze the CA data of Figure 3 within the Butler-Volmer formalism, one needs to present the current density data as function of the overpotential, rather than as a function of the applied electrode potential. The overall current density j is directly correlated with the produced chlorine as long as the CER constitutes the dominating reaction, i.e. OER and other parasitic side effects are negligible. This is accomplished at pH = 0.9 as discussed later in Figure 5. Secondly, an increase in temperature results in an exponential rise of the exchange current density j0, a measure of the intrinsic catalytic properties of an electrode material, which is mainly determined by the apparent free activation energy G#rds. In order to determine the apparent free activation barrier G#rds of CER, the experimentally measured current densities as a function of the applied electrode potential need to be transformed into Tafel plots (cf. Figure 4), which presents the overpotential ηCER as a function of the decade logarithm of the current density log (j). Typically for η > 30 mV the Tafel plot indicates a linear correlation between these two quantities in the so-called Tafel region. The exchange current density j0 for each of the four selected temperatures (table in Figure 4) are determined by extrapolating the corresponding Tafel lines to ηCER = 0 V. These data are required to derive the apparent free activation barrier of CER from an Arrhenius-like plot that will be discussed in Section 4 and critically compared to theoretical results. Similar data can be extracted from the OLEMS data and are shown in the supporting information (Figure S4).

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Figure 4: Presentation of the CA experiments in Figure 3 for temperatures of 25°C, 30°C, 35°C and 40°C in the form of Tafel plots. The extrapolation of the linear Tafel region (dotted lines) for each temperature to ηCER = 0 V provides the exchange current density j0 as a function of T presented in the inset table. In order to address the selectivity issue of the competing CER and OER, the OLEMS data are combined with the CA measurements by plotting the integrated mass signal m/z = 36 as function of the total charge transferred in the electrochemical reaction (Figure 5). It turns out that independent of the temperature all data points fall on a common line passing through the origin. The single line in Figure 5 is reconciled with Faraday’s law, which states that the mass of a single product is a linear function of the total charge that is independent of temperature. However, if two reactions (such as CER and OER) are competing with similar efficiencies in the considered potential range, for each temperature a separate line should occur (cf. Figure 7 in section 3.2). Therefore, we conclude from a single linear relation in Figure 5 that one reaction, namely CER, is dominating the electrolysis, which agrees well with industrial experience. In industry, about 1 – 3% of anodic gas production is oxygen at a pH of about 2. As our experiments were conducted at even lower pH value of about 0.9, the percentage of electrons entering any side reaction should be even lower so that the selectivity further shifts towards CER.

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Figure 5: OLEMS versus CA experiments for CER over RuO2(110) in a 5 M NaCl + 10−2 M HCl solution: The integrated m/z = 36 signal obtained by OLEMS is shown as function of the total charge transferred in the electrochemical reaction. 3.2. Experiments at pH = 3.5 The selectivity issue of the competing CER and OER was also studied at higher pH value by decreasing the concentration of hydrochloric acid to 10−5 M HCl, still retaining to a 5 M NaCl solution. The pH of the electrolyte solution was measured by a glass electrode and turned out to be 3.5.24 By increasing the pH value, the equilibrium potential of OER decreases, while that of the CER remains constant. Therefore, the overpotential of the OER becomes larger when the electrode potential is set to CER conditions, which in turn should be accompanied by a higher OER activity and therefore lower selectivity towards CER. CA and OLEMS experiments were conducted only for two temperatures, namely 25°C and 35°C (cf. Figure 6). Corresponding Tafel plots are displayed in the SI (Figure S5). Compared to Figure 3, the measured current densities at pH = 3.5 is by one order of magnitude larger than that at pH = 0.9, which may indicate that the OER or other electrochemical side reactions are dominating at pH = 3.5. The m/z = 36 signal in Figure 6 is, however, smaller compared to Figure 3. We have to bear in mind that between the measurements at pH = 0.9 and pH = 3.5 the electrolyte solution was exchanged and the distance of the OLEMS tip to the catalyst’s surface might be different. As a consequence, the OLEMS m/z = 36 signal cannot quantitatively be compared between the two pH values. Since the CER over RuO2(110) is independent of pH,23,27

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one would expect that the amount of gaseous chlorine detected by OLEMS remains unchanged to a first approximation.

Figure 6: CA and OLEMS experiments of a RuO2(110)/Ru(0001) working electrode in 10−5 M HCl (pH = 3.5) and 5 M NaCl subjected to a series of electrode potential pulses recorded at a) T = 25°C and b) T = 35°C in the CER regime. The selectivity issue at pH = 3.5 is again studied by plotting the integrated mass signal of m/z = 36 as a function of the total charge in the electrochemical reaction (cf. Figure 7). Quite in contrast to pH = 0.9 (cf. Figure 5), there is not a single line, but rather for each temperature we recognize a non-linear increase of the integrated mass signal as function of the total charge. This observation is not compatible with Faraday’s law for a single reaction (CER) but instead point towards the presence of pronounced competing reactions.

Figure 7: Integrated area of m/z = 36 ion current versus total charge measured by CA per potential pulse for CER over RuO2(110) in a 5 M NaCl + 10−5 M HCl solution: The integrated m/z = 36 signal obtained by OLEMS is shown as function of the total charge transferred in the electrochemical reaction. - 13 ACS Paragon Plus Environment

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4. Discussion 4.1. Selectivity The selectivity of the competing CER and OER over a single-crystalline RuO2(110) model electrode was studied in Section 3 by varying the pH value. It turns out that for pH = 0.9 (cf. Section 3.1), the electrons consumed are used for the production of gaseous chlorine rather than for the formation of gaseous oxygen. From a thermodynamic point of view, OER is preferred over CER according to the lower standard equilibrium potential (U0OER = 1.23 V vs. SHE, U0CER = 1.36 V vs. SHE). However, the kinetics of CER and OER differ significantly from each other, since CER is a two-electron process, whereas OER is a sluggish four-electron process. Therefore, OER requires substantially larger overpotentials than CER in order to maintain the same current density. In the operating potential window (1.26 V up to 1.40 V vs. SHE) CER constitutes the dominating reaction if the pH of the solution is kept low (pH = 0.9). This conclusion is supported most notably by Figures 4 and 5 in that the Tafel slope and the Faradaic efficiency are shown to be independent of temperature. Since OER and CER reveal different apparent symmetry factors,23,26,27 a change in the value of the Tafel slope with temperature would be expected when OER competes substantially with CER. Equally, CER and OER reveal different apparent free activation energies so that a substantial competition of CER and OER would cause a non-linear behavior in Figure 5, which is not observed. Altogether, we conclude that for pH = 0.9 the CER is dominating the potential region above 1.30 V vs. SHE and therefore the apparent free activation barrier of CER can unambiguously be determined from the temperature-dependent Tafel plots of Figure 4. In section 4.2 the apparent free activation barrier of CER over RuO2(110) is derived from temperature-dependent measurements at pH = 0.9 and is critically compared to the free energy landscape along the reaction coordinate as obtained from density functional theory (DFT) calculations.14 The selectivity situation changes dramatically when running the CER under higher pH values. For pH = 3.5 neither the Tafel slopes nor the Faradaic efficiency remains independent of the temperature. This is a clear indication of competing reactions at pH = 3.5. Comparison of the current density in Figure 3 and Figure 6 indicates that the CER is not dominating the current density at pH = 3.5. The competing reaction can be the OER, but we have to note that at higher pH value also the formation of hypochlorous acid and hypochlorite OCl− becomes significant.7 This will further lower Faraday efficiency. Unfortunately, the CER versus OER selectivity could not be quantified in the present study as the sensitivity of the MS towards molecular oxygen is - 14 ACS Paragon Plus Environment

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too low in the presence of large amounts of Cl2.22 In Figure S6 we show MS data of m/z = 32 (oxygen) and m/z=36 (chlorine) for pH = 0.9 and 3.5 at T = 35°C. The high m/z = 32 OLEMS signal is mainly due to water dissociation at the MS filament, while the small undulations in the m/z = 32 signal when applying a potential pulse are counteracting the chlorine evolution. The reason can either be traced to a lowering of the MS sensitivity to oxygen when chlorine is present in the chamber or by replacement of dissolved oxygen by dissolved chlorine in the electrolyte solution. Even without any oxygen evolution we can detect a substantial oxygen signal in OLEMS that is coming from water dissociation, thus generating a high background O2 level. Future work will be devoted to optimize the OLEMS system for detecting oxygen in the presence of chlorine.

4.2. Apparent free activation energy of CER: Microkinetic Modeling Recently, the reaction mechanism of CER over RuO2(110) was resolved by microkinetics from first principles.14 Essentially there are three reasonable reaction mechanism for the CER over RuO2 discussed in the literature, namely Volmer-Tafel,28 Volmer-Heyrovsky,29 and Krishtalik.30 A recent ab initio TD study31 favored the Krishtalik mechanism based on the adsorption of O2 on RuO2(110) which, however, was shown not to be stable against dissociation.32 From ab initio kinetics it turns out that the CER proceeds via a Volmer-Heyrovsky mechanism in which the adsorption and discharge of chloride (Volmer step) from the electrolyte solution on the fully oxygen-covered RuO2(110) surface32,33 is followed by the direct recombination of the adsorbed chlorine species (OClot) and another chloride anion from the electrolyte solution (Heyrovsky step): (i) Oot + 2Cl− → OClot + e− + Cl−

(Volmer step)

(ii) OClot + e− + Cl− → Oot + Cl2 + 2e−

(Heyrovsky step)

The corresponding free energy profile along the reaction coordinate for zero overpotential, i.e. forward and backward reaction are in thermodynamic equilibrium, is depicted in Figure 8. The free energy diagram reveals two transition states (TS), labeled as #1 and #2, and one reaction intermediate (RI), OClot. Under typical reaction conditions the reaction intermediate is in equilibrium with the reactant, establishing a so-called quasi-equilibrium. Therefore, the ratedetermining reaction (rds) step is identified with the TS being highest in free energy Grds# 34 and the reaction rate r per active site either in cathodic or in anodic direction can be expressed in the following manner: - 15 ACS Paragon Plus Environment

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#  −Grds  k BT r (η = 0V ) = ⋅ exp  (2)  h  k BT  In equation (2) h denotes the Planck constant. By using Faraday’s law, equation (2) translates to

the exchange current density j0 for η = 0 V: #  −Grds  k BT ⋅ zeΓ act j0 = ⋅ exp  (3)  h  k BT  Here, z and Γact are the number of electrons in the cell reaction (CER: z = 2) and the number of

active sites per area, which amounts to 5·1014 cm−2 for RuO2(110) according to surface oxygen Oot being the active center for CER.14 The use of a model electrode, RuO2(110), is decisive as otherwise Γact is unknown.

Figure 8: The free energy diagram of the CER over RuO2(110) for η = 0 V as determined by a recent full kinetics study from first principles based on the Volmer-Heyrovsky mechanism.14 The free energy landscape along the reaction coordinate (Figure 8) reveals that TS #2 is higher in free energy than TS #1 for η = 0 V. Assuming quasi-equilibrium the second reaction step (Heyrovsky step) is therefore identified with the rds for zero overpotential with an apparent free activation energy of Grds# = G2# = 0.79 eV. Apart from the theoretically obtained value of the apparent free activation energy, the temperature-dependent measurements (Figure 4) enable to determine Grds# by using the following equation, which is derived by taking the natural logarithm on both sides of equation (3): # k T ⋅ zeΓact  −Grds  ln  j0 (T) / B =  h   kBT

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In Figure 9 are also corresponding OLEMS data for m/z = 36 (cf. Figure S4) included. From the slope of the Arrhenius-like plot in Figure 9, the apparent free activation energy Grds# is derived to be 0.91 eV, which is in reasonably good agreement with the theoretical value of 0.79 eV. The temperature dependence of Γact is according to the performed DFT calculations negligibly small, since the adsorption of chlorine on the catalyst surface is strongly endergonic for η = 0 V (Figure 8, OClot is 0.34 eV higher in free energy than Oot) and therefore, almost all active sites are available for the adsorption and discharge of chlorine independent of temperature. Consequently, blocking effects can be clearly ruled out. The difference between the experimentally and theoretically derived apparent activation energy is 0.12 eV, which is within the confidence interval of current state of the art theoretical models considering in particular the approximate treatment of solvent effects.

Figure 9: The exchange current density j0 as derived from the temperature-dependent Tafel plots in Figure 4 and the OLEMS data from Figure S4 are converted in an Arrhenius-like plot. From the slope an apparent activation energy of 0.91 eV can be deduced. We may recall that in electrochemistry a variation of the electrode potential or the overpotential is always associated with changes in the free energy so that with temperature-dependent experiments we have only access to the apparent free activation energy Grds#. This is quite in contrast to heterogeneous catalysis with chemical reactions, where instead of Grds# the apparent activation energy is mostly determined by Arrhenius plots. - 17 ACS Paragon Plus Environment

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4.3 Extrapolation to industrial conditions In industrial chlor-alkali electrolysis, the CER is carried out at 80°C in a 5 M NaCl solution.4 Under these conditions, the typically applied overpotential of around 100 mV is accompanied by virulent Cl2 bubble formation. In Figure 10 the overpotentials, for which bubble formation sets in, are indicated as red dots dependent on the investigated temperatures of 25°C, 30°C, 35°C, and 40°C; the onset of bubble formation was visually detected by naked eye. In the SI, we show that the onset overpotential for the formation of bubbles is a linear function of temperature, which enables to extrapolate the straight line given by the four experimental points to higher temperatures. It turns out that under industrial conditions (80°C) the onset potential for bubble formation amounts to 82 mV. This value agrees well with the applied overpotential of 100 mV where industrial CER is operating under virulent bubble formation.

Figure 10: Threshold overpotentials for bubble formation (red) as a function of temperature extrapolated to industrial conditions (80°C) amounts to 82 mV (green). This value is in fairly good agreement with the operating conditions of the industrial CER process. In addition, the sensitivity of OLEMS for m/z = 36 is shown (blue). With OLEMS, we were able to monitor evolving chlorine at much lower overpotentials. In addition, we indicate in Figure 10, at which onset overpotential OLEMS is able to detect evolving chlorine. CER can be detected with OLEMS at much lower overpotential when no bubbles have been formed yet. This higher sensitivity is an obvious advantage of capillary-based - 18 ACS Paragon Plus Environment

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OLEMS over standard, large area detection DEMS.18,35 However, as soon as bubbles appear, the OLEMS signal fluctuates severely whenever a bubble is in front of the OLEMS tip, so that OLEMS practically cannot be used for high overpotentials. In this regime, however, DEMS could provide a more reliable, quantifiable signal.

5. Conclusions Ultrathin single crystalline transition metal oxide films, in our case RuO2(110) supported on Ru(0001), are prepared under ultra-high vacuum conditions serving as model electrodes in an electrocatalyzed reaction under well-defined electrochemical conditions. Only such model studies allow for a tight connection between experiment and theory and the prospect to improving our molecular understanding of electrocatalyzed reactions.36 In particular, we studied the activity and selectivity of the chlorine evolution reaction (CER) over RuO2(110) at pH = 0.9 and pH = 3.5, employing the techniques of chronoamperometry (CA) and on-line electrochemical mass spectrometry (OLEMS) at various reaction temperatures. As shown by combined OLEMS/CA experiments (cf. Figure 5), the OER versus CER selectivity shifts from pure CER at pH = 0.9 towards OER and other parasitic reactions at pH = 3.5 (cf. Figure 7). The apparent free activation energy is derived from the temperature dependence of the exchange current density j0 of the CER over RuO2(110) (cf. Figure 9). In order to extract j0(T) from temperature-dependent current density versus electrode potential data, we adjust the reaction conditions in such a way that practically only the CER takes place (pH = 0.9) and that the temperature dependence of the ion transport (U < 1.4 V) can be ignored, and we properly account for the temperature variation of the reversible half-cell potential UCER. The experimentally derived value of apparent free activation energy of 0.91 eV can be directly compared to the free energy diagram along the reaction coordinate (cf. Figure 8) that was recently determined by first principles calculations.14 Assuming quasi-equilibrium for the microkinetic modeling an apparent free activation energy of 0.79 eV is inferred. We consider the experimentally determined apparent free activation energy of 0.91 eV as a benchmark, against which future improved theoretical modeling from first principles can be assessed.

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Temperature dependence of the reference electrode, typical mass spectrum taken by OLEMS, determination of the equilibrium potential of the CER over RuO2(110), pH dependence of the CER over RuO2(110), Tafel plots of the OLEMS data for pH = 0.9, Tafel plots for pH = 3.5 , OLEMS experiments including the m/z = 32 MS signal, derivation of the linear temperature behavior for the onset of bubble formation. This material is available free of charge via internet at http:///pubs.acs.org.

Acknowledgement: The authors like to thank Ad Wonders of Eindhoven University of Technology for support with the OLEMS measurements and fruitful discussions on the technical aspects of OLEMS. ISE thanks Prof. Baltruschat, University Bonn for his kind introduction into surface electrochemistry. We thank Dr. Benjamin Herd and Dr. Philipp Krause from JLU for providing us the LEED data in Figure 1, while we thank Tim Weber for the measurements of the pH dependence of the CER in the SI, Figure S3. AG and EJMH acknowledge funding by the Dutch National Research School Combination Catalysis Controlled by Chemical Design (NRSC-Catalysis). EJMH acknowledges support from an NWO TOP grant. HO thanks financial support from BMBF (HEXCHEM).

References (1) Schmittinger, P.; Florkiewicz, T.; Curlin, L. C.; Lüke, B.; Scannell, R.; Navin, T.; Zelfel, E.; Bartsch, R. In Ullmann’s Encyclopedia of Industrial Chemistry, Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim, Germany, 2012. (2) Karlsson, R.K.B.; Cornell, A. Chem. Rev. 2016, 116, 2982-3028. (3) a) Eurochlor, Chlorine Industry Review 2013−2014. b) http://eippcb.jrc.ec.europa.eu/reference/BREF/CAK_Adopted_072014.pdf (4) Moussallem, I.; Jörissen, J.; Kunz, U.; Pinnow, S.; Turek, T. J. Appl. Electrochem. 2008, 38, 1177-1194. (5) Gestermann, F.; Ottaviani, A. In Modern Chlor-Alkali Technology; Moorhouse, J., Eds.; Blackwell Science, London, 2001, Vol. 8, p 49-56. (6) Dionigi, F.; Reier, T.; Pwolek, Z.; Gliech, M.; Strasser, P. Chem. Sus. Chem. 2016, 9, 962972. (7) Zeradjanin, A.R.; Menzel, N.; Schumann, W.; Strasser, P. Phys. Chem. Chem. Phys. 2014, 16, 13741-13747. (8) Beer, B. H. U.S. Patent 3096272, 1963. (9) Beer, H.B. J. Electrochem. Soc. 1980, 127, 303C-307C. (10) Hayfield, P. C. S. Platinum Met. Rev. 1998, 42, 27−33. - 20 ACS Paragon Plus Environment

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(11) Trasatti, S. Electrochim. Acta 2000, 45, 2377−2385. (12) Zeradjanin, A.R.; Schilling, T.; Seisel, S.; Bron, M.; Schuhmann, W. Anal. Chem. 2011, 83, 7645. (13) Herd, B.; Knapp, M.; Over, H. J. Phys. Chem C. 2012, 116, 24649-24660. (14) Exner, K.S.; Anton, J.; Jacob, T.; Over, H. Angew. Chem. Int. Ed. 2016, 55, 7501-7504. (15) Exner, K.S.; Anton, J.; Jacob, T.; Over, H. Electrocatal. 2015, 6, 163. (16) Markovic, N.M.; Grgur, B.N.; Ross, P.N. J. Phys. Chem. B 1997, 101, 5405-5413. (17) Perales-Rondon, J.V.; Herrero, E.; Feliu, J.M. J. Electroanal. Chem. 2015, 742, 90-96. (18) Wonders, A.H.; Housmans, T.H.M.; Rosca, V.; Koper, M.T.M. J. Appl. Electrochem. 2006, 36, 1215-1221. (19) Bruckenstein, S.; Gadde, R. R., J. Am. Chem. Soc. 1971, 93, 793–794. (20) Wolter, O.; Heitbaum, J. Ber. Bunsenges. 1984, 88, 2—6. (21) Baltruschat, H. J. Am. Soc. Mass. Spectrom. 2004, 15, 1693—1706. (22) Hammes, M.; Valtchev, M.; Roth, M.B.; Stöwe, K.; Maier, W.F. Appl. Catal. B: Environmental 2013, 132-133, 389. (23) Guerrini, E.; Consonni, V.; Trasatti, S. J Solid State Electrochem. 2005, 9, 320-329. (24) The co-addition of 5 M NaCl in a HCl solution decreases the pH value significantly: For 102 M and 10-5 M HCl the pH values measured with a glass electrode amount to 1.95 and 4.98 (without NaCl) or 0.9 and 3.5 (with 5 M NaCl), respectively. (25) Bockriss, J. O’M.; Ready, A.K.N. In Modern Chemistry, A Plenum/Rosetta Publication, 1973, Vol. 2. (26) Castelli, P.; Trasatti, S.; Pollak, F.H.; O’Grady, W.E. J. Electroanal. Chem. 1986, 210, 189194. (27) Consonni, V.; Trasatti, S.; Pollak, F.H.; O’Grady, W.E. J. Electroanal. Chem. 1987, 228, 393-406. (28) Trasatti, S.; O’Grady, W. E. In Advances in Electrochemical Science and Engineering, Gericher H., Tobias, C.W., Eds.; Wiley, New York, 1981, Vol. 12, p 117-128. (29) Jansen, L. J. J.; Starmans, L. M. C.; Visser, J. G.; Barendrecht, E. Electrochim. Acta 1977, 22, 1093-1100. (30) Krishtalik, L. I. Electrochim. Acta, 1981, 26, 329-337. (31) Hansen, H.A.; Man, I.C.; Studt, F.; Abild-Pedersen,F.; Bligaard, T.; Rossmeisl, J. Phys. Chem. Chem. Phys. 2010, 12, 283-290. (32) Exner, K.S.; Anton, J.; Jacob, T.; Over, H. Electrochim. Acta 2014, 120, 460-466. (33) Exner, K.S.; Anton, J.; Jacob, T.; Over, H. Angew. Chem. Int. Ed. 2014, 53, 11032-11035. (34) Parsons, R. Trans. Faraday Disc. 1951, 47, 1332-1334. (35) Jusys, Z.; Massong, H.; Baltruschat, H. J. Electrochem. Soc. 1999, 146, 1093-1098. (36) Markovic, N.M.; Ross Jr., P.N. Surf. Sci. Rep. 2002, 45, 117-229.

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