Selectivity between Oxygen and Chlorine Evolution in the Chlor-Alkali

Review pubs.acs.org/CR

Selectivity between Oxygen and Chlorine Evolution in the ChlorAlkali and Chlorate Processes Rasmus K. B. Karlsson and Ann Cornell* Applied Electrochemistry, School of Chemical Science and Engineering, KTH Royal Institute of Technology, SE-100 44 Stockholm, Sweden ABSTRACT: Chlorine gas and sodium chlorate are two base chemicals produced through electrolysis of sodium chloride brine which find uses in many areas of industrial chemistry. Although the industrial production of these chemicals started over 100 years ago, there are still factors that limit the energy efficiencies of the processes. This review focuses on the unwanted production of oxygen gas, which decreases the charge yield by up to 5%. Understanding the factors that control the rate of oxygen production requires understanding of both chemical reactions occurring in the electrolyte, as well as surface reactions occurring on the anodes. The dominant anode material used in chlorate and chlor-alkali production is the dimensionally stable anode (DSA), Ti coated by a mixed oxide of RuO2 and TiO2. Although the selectivity for chlorine evolution on DSA is high, the fundamental reasons for this high selectivity are just now becoming elucidated. This review summarizes the research, since the early 1900s until today, concerning the selectivity between chlorine and oxygen evolution in chlorate and chlor-alkali production. It covers experimental as well as theoretical studies and highlights the relationships between process conditions, electrolyte composition, the material properties of the anode, and the selectivity for oxygen formation.

CONTENTS 1. Introduction 1.1. Industrial Chlor-Alkali Production 1.2. Industrial Sodium Chlorate Production 1.3. Selectivity Concept in Chlor-Alkali and Sodium Chlorate Production 2. Selectivity for Oxygen and Chlorine Evolution in Chlor-Alkali and Chlorate Processes 2.1. Studies on Oxygen Formation during Decomposition of Hypochlorous Acid Species 2.1.1. Uncatalyzed Decomposition 2.1.2. Catalyzed Decomposition 2.2. Studies on Selectivity during Electrolysis 2.2.1. Academic Research from 1900 to 1969 2.2.2. Academic Research during the 1970s 2.2.3. Research during the 1980s 2.2.4. Research during the 1990s 2.2.5. Research during the 2000s 2.3. Electrochemical Studies Concerning the Relationship between Selectivity and Stability 3. Discussion 3.1. Factors Affecting the Selectivity between Chlorine and Oxygen Evolution in the Production of Chlorine and Chlorate 3.1.1. Influence of Process Conditions 3.1.2. Influence of Additions and Contaminants in the Electrolyte 3.1.3. Influence of the Anode Structure and Composition © 2016 American Chemical Society

3.2. Extent of Oxygen Evolution from Different Reactions 3.3. Connection between Anode Stability and Selectivity 4. Conclusions Author Information Corresponding Author Notes Biographies Acknowledgments Nomenclature References

2982 2983 2985 2986 2986 2986 2986 2987 2988 2988 2990 2991 2997 3000

3019 3020 3021 3021 3021 3021 3021 3022 3022 3022

1. INTRODUCTION Chlorine gas and sodium chlorate are base chemicals, with yearly production rates of 50−60 million tonnes1,2 and over three million tonnes,3 respectively. Chlorine gas is used in a wide variety of applications, including in production of construction materials such as polyvinyl chloride (PVC), in organic synthesis, in metallurgy, and in water treatment.4 Sodium chlorate finds its use almost exclusively for bleaching of pulp and paper products.3 Both products are produced predominantly through electrolysis of sodium chloride brine. Although the charge yields of both the chlor-alkali process, in which chlorine gas and sodium hydroxide are produced, and chlorate processes are quite high, ≥ 95%,5,6 the fact that both chlorine gas and chlorate are bulk chemicals,

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Received: July 6, 2015 Published: February 16, 2016

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DOI: 10.1021/acs.chemrev.5b00389 Chem. Rev. 2016, 116, 2982−3028

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ature, affect reactions both in the bulk electrolyte and on the anode surface. This review therefore discusses all three aspects: electrolyte composition, anode composition, and the process conditions. The focus of this review is on the selectivity issue in general, without limitation with regard to the anode material used. Since there are connections between some mechanisms of electrode decomposition and oxygen production, the stability of DSA will also be discussed. The following section of the introduction constitutes a brief overview of the chlor-alkali and chlorate processes, as well as the most important and well-known main and side reactions. It gives the necessary basic information to describe effects of various process, electrolyte, and anode properties on the selectivity. The selectivity in the chlor-alkali and chlorate processes is not a simple concept, as it involves both direct formation of oxygen on the anode as well as oxygen resulting from bulk reactions. Therefore, the selectivity concept as applied for industrial chloralkali and sodium chlorate production, as well as some of the methods that have been used to explore it, will be described in section 1.3. Afterward, in section 2, the literature relevant to the selectivity is described, in chronological order from the early 1900s until today. The literature is described in three main sections, covering research on oxygen-evolving reactions relevant to the electrolyte processes, electrochemical studies (both experimental and theoretical), and studies on the link between stability and selectivity. Finally, a discussion of the results and trends found in literature is given, as well as some conclusions regarding future research needs.

produced in large quantities, means that losses of even a few percent are costly. The yearly electricity consumption in chloralkali production alone is about 150TWh,1 which makes it clear that improvements in energy efficiency of these processes could yield significant reductions in world energy consumption. The main side reaction in industrial electrolytic cells is the production of oxygen. In the chlorate process, the unwanted oxygen combines with hydrogen in the cell gas and may form explosive gas mixtures, which also makes oxygen a safety hazard. Further study of the factors deciding the selectivity between chlorine or chlorate and oxygen production in industrial chlor-alkali and chlorate cells is therefore warranted. However, there is no review of the scientific literature that focuses specifically on the selectivity issue. This review aims to bridge that gap. The electrolytic production of chlor-alkali and sodium chlorate is performed in either divided, for production of chlor-alkali, or undivided, for production of sodium chlorate, electrochemical cells using, for example, steel or activated Ni cathodes (Ni cathodes are only used in chlor-alkali production) and dimensionally stable anodes (DSA, dimensionally stable anode, is a trademark of Industrie De Nora S.p.A., Italy.).1,7 DSA are mixed ruthenium−titanium oxide (RTO) coatings of rutile RuO2 and TiO2 deposited on Ti. The usage of DSA to describe anodes in the present review is primarily used to refer to anodes of the typical industrial “Beer 2” (30% RuO2 and 70% TiO2 coated on Ti) formulation.8 Commercial DSA electrodes most often contain one or more other additional dopant materials. RTO is used in the present review as a more general term for anodes with coatings containing varying amounts of RuO2 and TiO2. In these coatings, the RuO2 and TiO2 components form solid solutions, where RuIV and TiIV are part of the same rutile lattice.9−11 While rutile RuO2 has a high electronic conductivity,12 pure rutile TiO2 is a semiconductor with a band gap of 3 eV.13 Nevertheless, the mixed oxide has a high electronic conductivity, enabling its use as an electrode, as the doping with RuIV introduces new electronic states in the region of the TiO2 band gap.14,15 The preparation and usage of dimensionally stable anodes was patented by Beer in a series of patents in the 1960s (Britain) and 1970s (United States).8,16−23 The discovery of DSA has been called “one of the greatest technological breakthroughs of the past 50 years of electrochemistry”.24 Since then, the usage of DSAs in these processes has led to significant energy savings due to their lower potentials at industrial current densities.24,25 However, as the name implies, their most important advantage over previous graphite electrodes is their stability, with modern DSAs being able to operate at industrial current densities for more than 10 years.24 The literature concerning DSAs and RTOs in general is summarized in several reviews.23−33 The recent reviews of Over concerning the chemistry and catalysis of pure RuO212 and concerning correlations and contrasts between liquid phase electrochemical and gas phase chemical reactions on RuO2 and related metallic oxide materials,34 as well as the review of Diebold35 concerning the surface chemistry of TiO2, should also be mentioned. The oxygen evolution side reaction (OER) in chlor-alkali and sodium chlorate production is connected to catalytic processes ongoing both in the electrolyte and on the anode surface, where oxygen might be evolved both electrochemically and chemically through water or hypochlorous acid (or hypochlorite) decomposition. Therefore, to get a complete picture of the selectivity issue in these industrial processes, compositions of both the electrolyte and the anode are important. Furthermore, other process conditions, such as current density and temper-

1.1. Industrial Chlor-Alkali Production

Chlorine gas and sodium hydroxide are produced in chlor-alkali production according to the following overall reaction1 2NaCl + 2H 2O → Cl 2 + H 2 + 2NaOH

(1)

The electrochemical formation of chlorine gas at the anode (DSA) occurs according to the following reaction 2Cl− → Cl 2 + 2e−

(2)

with an associated equilibrium potential of E° = 1.3583 V versus SHE, where the standard potentials used are those of Bard et al.,37 unless otherwise indicated. The evolved chlorine gas is collected and processed in a number of steps. As chlorine is being evolved at the anode, hydrogen (except in the mercury cell process) forms at the cathode 2H 2O + 2e− → H 2 + 2OH−

(3)

a reaction with E° = −0.8277 V. The anode compartment is kept separated from the cathode compartment, and the way in which they are separated constitutes the main characteristic of the different processes used to produce chlor-alkali. The mercury cell process, the diaphragm process, and the membrane process dominate industrial production of chlor-alkali. In a membrane cell, a cation-selective membrane separates the anode and cathode compartments, whereas a microporous diaphragm is used in the diaphragm process. The first industrial process was the mercury cell process, in which a mercury amalgam is produced at the cathode instead of hydrogen. The amalgam is then later decomposed in a separate reactor to yield hydrogen gas and sodium hydroxide. Today, the mercury cell and diaphragm processes are being replaced by the membrane process for both environmental and technical reasons.38,39 By the end of year 2013, the total installed chlorine capacity in Europe consisted of 59% membrane technology, 25% mercury cells, and 14% 2983


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diaphragm cells.39 The usage of mercury cells in the European Union is scheduled to be discontinued in 2017.39 Typical conditions applied in the membrane process are shown in Table 1, and more detailed descriptions of the three technologies can be

Table 2. Some Important Dissolved and Gaseous Chlorine Species and Their Cl Oxidation States.42

Table 1. Typical Conditions for Chlor-Alkali Production in Membrane Processes1,2,36 cell voltage (V) current density (kA m−2) temperature (°C) NaCl concentration in the anolyte (g dm−3) anolyte pH NaOH concentration in the catholyte (wt %)

2.4−2.7 1.5−7 90 200 2−4 32

name

Cl oxidation state

chloride (ion) trichloride ion hydrochloric acid chlorine gas hypochlorous acid [or chloric (I) acid] hypochlorite (ion) [or chloric (I) ion] chlorine monoxide chlorous acid chlorite (ion) chlorine dioxide chlorate (ion) perchlorate (ion)

−1 −1, 0 −1 0 +1

ClO− (aq) Cl2O (g) HClO2 (aq) ClO−2 (aq) ClO2 (g) ClO−3 (aq) ClO−4 (aq)

found in, for example, Wendt et al.40 or Schmittinger et al.1 A small percentage of the total chlorine production is based on electrolysis of 17 wt % HCl in either diaphragm or membrane cells.1 The energy cost for the production is usually given in terms of kWh t−1 NaOH produced, with 886 kg Cl2 being produced per tonne of NaOH. In membrane cells, the specific electrical energy requirement is approximately 2100 kWh t−1 NaOH.1 Dissolved chlorine may hydrolyze according to the dynamic equilibrium in reaction 4. Formed hypochlorous acid is involved in another equilibrium with water, reaction 5. The impact of the pH value on the relative concentrations of Cl2, HOCl, and ClO− is significant, as is indicated in Figure 1. Cl 2 + H 2O ⇌ H+ + HOCl + Cl−

(4)

HOCl + H 2O ⇌ ClO− + H3O+

(5)

Hypochlorous acid and hypochlorite react to form chlorate, ClO−3 , according to reaction 6. 2HOCl + ClO− → ClO−3 + 2H+ + 2Cl−

species Cl− (aq) Cl−3 (aq) HCl (g) or (aq) Cl2 (g) or (aq) HOCl (aq)

+1 +1 +3 +3 +4 +5 +7

Figure 2. Pourbaix diagram for aqueous species in the chlorine-water system, indicating the stable equilibria at 25 °C and unit activity of the dissolved species. Figure prepared using pymatgen45−47 and thermodynamic data from Wagman et al.48

(6)

Figure 1. Percentages of active chlorine species Cl2, HOCl, and ClO− as a function of anolyte pH at 90 °C and 200 g NaCl dm−3. The percentages are calculated using equilibrium constants for Cl2 hydrolysis and HOCl deprotonation. Reprinted with permission from ref 41. Copyright 1990 Springer Science and Business Media.

Figure 3. Pourbaix diagram for the aqueous metastable chlorates (excluding ClO−4 ) in the chlorine-water system, indicating the equilibria at 25 °C and unit activity of the dissolved species. Figure prepared using pymatgen45−47 and thermodynamic data from Wagman et al.48

Cl−, HOCl, ClO−, and ClO−3 are the main species of interest during chlor-alkali and chlorate production, but the chlorinewater equilibrium in aqueous solution is significantly more complex. Table 2 shows some of the known dissolved and gaseous species that might occur during electrolysis.42,43 The stability of these species at different electrode potentials and pH values is indicated in Figures 2 (stable species), 3 (metastable chlorate species), 4 (metastable chlorite species), and 5 (metastable hypochlorite species). The trichloride ion, Cl−3 ,

which is in equilibrium with Cl2 (aq) under chlor-alkali conditions (chloride activities between 2 and 6),44 is not seen in the figures, as its window of stability is very small. In addition to chlorate ions, hypochlorite (or hypochlorous acid) may also form oxygen as will be described in the next section. Both chlorate ions and oxygen gas are highly undesired in chlor-alkali cells, and thus, the formation of the reactant hypochlorite species should be suppressed. In the membrane and 2984


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NaCl + 3H 2O → NaClO3 + 3H 2

(7)

The energy requirement in kWh per tonne sodium chlorate produced in the industrial process can be estimated using (with a typical cell voltage of 3.1 V)51 Pe = 1511 Ah t−1 × 3.1V ≈ 4700 kWh t−1

(8)

The same electrode reactions that occur in the chlor-alkali process, reactions 2 and 3, are also believed to be the first steps in chlorate production.51 Dimensionally stable anodes and cathodes of either low carbon steel or titanium are used in the electrolysis cell.6 The chlorate cells are undivided, and the bulk electrolyte pH is 6−7. However, the ongoing anode and cathode reactions make the pH at the anode more acidic, while the pH at the cathode is alkaline. Thus, a pH gradient is set up across the cell, and this means that the hydroxide ions that form in the cathode reaction 3 can hydrolyze the chlorine that forms in the anode reaction 2.6,7,51,52 This gives the following dynamic equilibrium

Figure 4. Pourbaix diagram for the aqueous metastable chlorites (excluding ClO−3 and ClO−4 ) in the chlorine-water system, indicating the equilibria at 25 °C and unit activity of the dissolved species. Figure prepared using pymatgen45−47 and thermodynamic data from Wagman et al.48

Cl 2 + 2OH− ⇌ ClO− + H 2O + Cl−

(9)

Alternatively, water can hydrolyze the chlorine gas according to reaction 4.6 Hypochlorite is formed through reactions 4 and 5, or through reaction 9, and then chlorate is formed according to reaction 6.51 This reaction is a purely homogeneous chemical reaction, which has a maximum rate around pH 6.1−6.4 under industrial conditions and 80 to 90 °C.7 The pH that is used in the chlorate process is chosen to prevent chlorine gas formation, which increases in rate at lower pH, and oxygen gas formation, which increases in rate at higher pH. Chlorate cells usually have a large volume tank through which the reaction solution flows to allow reaction 6 to reach high conversion.6,40,53 The process temperature is kept high to increase the rate of reaction 6 (enabling the usage of a smaller reaction vessel) and of the electrode reactions (lowering the cell voltage).6,53−56 However, the rate of parasitic oxygen evolution also increases with increasing temperature, meaning that the temperature chosen is a balance between minimizing the rate of oxygen evolution and maximizing the reaction rates of the desired reactions.57 A number of anodic reactions limit the charge yield, particularly for chlorate production but also for chlor-alkali production. One of the more well-known is the following anodic side-reaction in which both oxygen and chlorate are formed:58,59

Figure 5. Pourbaix diagram for the aqueous metastable hypochlorites (excluding ClO−3 , ClO−4 , HClO2, and ClO−2 ) in the chlorine-water system, indicating the equilibria at 25 °C and unit activity of the dissolved species. Figure prepared using pymatgen45−47 and thermodynamic data from Wagman et al.48

diaphragm cells, small amounts of hydroxide ions may pass through the separator, dividing the acidic anolyte and the alkaline catholyte, and increase the pH close to the separator in the anode chamber, with subsequent oxygen and chlorate generation. Careful acidification of the anolyte can counteract the increase in pH related to the hydroxide leakage, resulting in oxygen levels in the chlorine gas of E > EO2/H2O), only the following two reactions were assumed likely to occur on the electrode: 2H 2O → O2 + 4H+ + 4e−

(21)

2Cl 2 + 4e− → 4Cl−

(22)

As these two reactions occur, 1 mol of gaseous oxygen and 4 mol of HCl forms. Thus, titration of the electrolyte was performed to measure the change in concentration of HCl a certain time after the RTO electrode was immersed in the solution. However, it was found that the change in pH corresponded directly to the amount of Cl2 that had been decomposed and that only a very small amount of ClO−3 had formed. The change in pH did not support spontaneous evolution of O2 under these conditions. This contrasts with the results of Kuhn and Mortimer,136 who noted oxygen evolution from NaCl solutions even without polarization, when RuO2 electrodes were used. Kokoulina and Bunakova then applied two methods to study the oxygen evolution under electrode polarization.155 First, titration was used to determine the amount of HCl formed during electrolysis, using this amount to then calculate the amount of oxygen formed. It was concluded that the oxygen evolution rate decreased with increasing Cl− concentration. The second method used potential compensation to obtain jO2, the partial current density for oxygen evolution, at different electrode 2993


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Figure 9. Voltammograms of DSA measured at increasing temperature from Kotowski and Busse.159 Reprinted from ref 159. Copyright 1986 John Wiley & Sons Ltd. All Rights Reserved. The figure may not be reproduced, stored in a retrieval system, or transmitted, in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, without the permission of Ellis Horwood Limited, Market Cross House, Cooper Street, Chichester, West Sussex, England.

In the same year, Buné et al.157 published results on the effect of the electrolyte composition on the kinetics and selectivity for OER and ClER on RTO anodes of the typical 30% RuO2-70% TiO2 composition. The measurements were carried out at a current density of 2000 kA/m2 in either 4.27 or 0.71 M NaCl, at 87 °C and pH 1.6. Effects of addition of either phosphate, sulfate, or perchlorate to the electrolyte were reported, continuing the work of Bondar et al.138 and Kazarinov and Andreev.143 After addition of phosphate, even at low concentrations of 0.004 to 0.1 M NaH2PO4, significant increases in anode potential were noted. The addition of 0.1 M of phosphate resulted in an increase in the selectivity for oxygen evolution, as measured by GC. Smaller effects were noted upon addition of similar amounts of sulfate, and no effect on the anode potential was noted upon addition of perchlorate. It was therefore concluded that phosphates and sulfates adsorb competitively, limiting the number of electrode surface sites accessible for chloride. The addition of increasing concentrations of phosphates also increased the electrode potential during oxygen evolution from solutions containing 4.27 M NaClO4, indicating the effect of competitive adsorption also on the oxygen evolution reaction. In 1986, Cairns et al.158 presented further work on ICI anode coatings for use in chlor-alkali membrane cells. It was noted that the “classical” method to prevent oxygen evolution was brine acidification, while modification of anodes was less frequently used. However, data for a modified “low oxygen” coating was presented, although without any details regarding the composition of this anode. It is probable that the coating was similar to those discussed earlier by Denton et al.154 In the same year, Kotowski and Busse159 presented research on the oxygen evolution side reaction in membrane cells. The results were based on measurements in both a diaphragm cell and in a membrane flow cell. Kotowski and Busse discussed the relative influence of reactions 10 (anodic chlorate production), 11 (water oxidation), 12 (hydroxide oxidation), 14 (bulk decomposition of hypochlorite), and also the reaction (first proposed by Wranglén160)

potentials. The method is based on the assumption that if the electrolyte is continuously saturated with Cl2-gas at a certain partial pressure, the electrode potential can be maintained at the reversible potential versus a RHE in the same solution. However, electrochemical oxygen evolution will cause the electrode potential to change from the reversible value. If a current is applied to maintain the electrode potential at Eanode = E0,Cl2, the current that is applied will be the same as that used for the side reaction. Under such conditions, the OER current is compensated by an equal and opposite current. To measure the OER current at different electrode potentials, the chlorine pressure of the gas above the electrolyte was decreased to lower and lower values, decreasing the equilibrium chlorine potential. The current that is applied is said to be the jO2(Eanode), yielding a polarization curve for oxygen evolution. The results showed 40 mV lower overpotentials for oxygen evolution at RuO2 than on RuO2+TiO2, indicating an increase in Cl2 selectivity for the mixed oxide. Effects on the overpotential of changes in pH were also noted. The results from both the titration and the potential compensation methods compared favorably. Also in 1984, Trasatti156 reported results of several studies of the oxygen and chlorine evolution on oxide materials, primarily pure oxides such as RuO2 or Co3O4. The bulk of the article discussed the interplay between surface area (q*) and point of zero charge in determining the activity of oxides in the relevant reactions. In the concluding part of the article, Trasatti correlated oxygen evolution overpotentials and chlorine evolution overpotentials for several oxide materials. He found that the slope between the two was close to unity. On materials, such as RuO2, where the overpotential for chlorine evolution is low, the overpotential for oxygen evolution is also low. He thought that this indicated that the selectivity between chlorine and oxygen evolution largely was not determined by the oxide material itself. While the activity for both oxygen evolution and chlorine evolution shows differences between different materials, the selectivity does not. 2994


Chemical Reviews HOCl + H 2O → 3H+ + Cl− + O2 + 2e−

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Based on these results, Kotowski and Busse159 concluded that in industrial cells, oxygen is evolved in reactions 10, 11, and 23. Two of the three reactions occur when hypochlorite reaches the anode. All three reactions involve water, leading the authors to suggest that the oxygen evolution overpotential measured in sulfuric acid could be used to predict the oxygen selectivity of the anode materials. The idea was supported by measurements on RTO anodes with decreasing amounts of RuO2 in the coating. For these electrodes, a linear relation between the logarithm of the oxygen content in the membrane cell gas and the overpotential for oxygen evolution, measured in sulfuric acid solution, was found (see Figure 10).

(23)

which is a combination of HOCl reduction (E° = 1.482 V) and H2O oxidation, making it relevant in the potential range 1.23 to 1.48 V versus SHE. Reactions 10 and 11 were found to make up between 20 and 60% of the total oxygen evolved in a membrane cell. Reaction 12 was not seen as important, since the pH at the anode surface is likely quite acidic when the electrode is anodically polarized. The amount of oxygen evolved due to reaction 11 was estimated based on data from amalgam cells, in which very little hypochlorite is formed. The amount of oxygen evolved due to reaction 10 was estimated based on the amount of chlorate that formed and the assumption that all chlorate formed in the membrane cell was formed due to oxidation of hypochlorite at the anode. Based on these initial experiments, Kotowski and Busse reached the conclusion that a large percentage (between 40 and 80%) of the total oxygen formed had to be formed in reactions involving hypochlorite, reactions 14 and 23. All other reactants for oxygen formation, such as Cl2O, ClO−2 , or HClO2, were excluded based on spectrophotometric data, leaving only hypochlorite species as reactants for oxygen evolution. Voltammograms of DSA were then measured in 200 g/dm3 NaCl and 3 to 5 g/dm3 NaClO. The voltammograms exhibited peaks at 400 to 500 mV and around 1000 mV (see Figure 9), with increasing intensity at increasing temperatures, which led to the hypothesis of a coupled redox reaction occurring at the anode. Kotowski and Busse thought it likely that concurrent oxidation, according to reaction 23, and reduction of hypochlorite 2e− + H+ + HOCl ⇌ Cl− + H 2O

(24)

occurred on the anode. The net result would then be production of oxygen according to 2HOCl → O2 + 2H+ + 2Cl−

(25)

To quantitatively test this hypothesis, further experiments were then performed at room temperature and high pH (pH ≈ 10, to prevent chlorine evolution), at which only reactions 10 and 23 were thought to occur. The electrolyte contained 30 g/dm3 NaCl and approximately 5 g/dm3 NaClO. Since the number of electrons released per mole of O2 formed is different for the two reactions (2 electrons for reaction 23 and 4 electrons for reaction 10), the percentage of the charge due to each of the reactions could be calculated. Three different anode coatings were used, a “normal”, a “special” (only said to be containing another platinum-group metal together with Ru and Ti), and a Pt/Ir coating. In industrial membrane cells, the gas evolved from the cell using the normal coating contained 0.45 vol % oxygen, the special coating gave a gas containing 0.17 vol % oxygen, and the Pt/Ir coating gave a gas containing 0.09 vol % oxygen. For the standard coating, the percentage of the oxygen evolved due to reaction 10 constituted 21% of the overall oxygen formation. For the special and the Pt/Ir coatings, the percentages of oxygen generated from the same reaction were 0% and 69%, respectively. In each case, the rest of the oxygen was accounted for by reaction 23, the coupled hypochlorite reduction/water oxidation reaction. Reaction 14, homogeneous decomposition of hypochlorite, was not considered to be important. Kotowski and Busse159 had done previous experiments that confirmed that several heavy metal ions promote the rate of this reaction. However, they did not find that solids, such as DSA or solid Pt/Ir, increased the rate of reaction.

Figure 10. Relationship between the overpotential for the OER and the volume percent of O2 in the off-gas, as found by Kotowski and Busse.159 The slope of the relationship between ηOER and log %O2 was found to be −169 mV per decade. Reprinted from ref 159. Copyright 1986 John Wiley & Sons Ltd. All Rights Reserved. The figure may not be reproduced, stored in a retrieval system, or transmitted, in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, without the permission of Ellis Horwood Limited, Market Cross House, Cooper Street, Chichester, West Sussex, England.

Again in 1986, Buné et al.68 published results using GC to measure the selectivity for oxygen formation using DSA. This time, the composition of the anode was changed; previous selectivity studies66,67 had examined the effect of current density, chloride concentration, and pH. The molar percentage of Ru in the coating was varied from 10% to 100%, similar to what was done in the study of Kotowski and Busse. The experiments were done in a similar way as in previously described articles by the same group.66,67 The conclusion from the measurements was that the selectivity for chlorine formation increased as the amount of Ru in the coating decreased, especially as the amount of Ru was decreased below 30%. The rationale was similar to that of Pecherskii et al.,67 that oxygen evolution is dependent on adjacent Ru sites, while chlorine evolution is not. Their results are summarized in Figure 11. 2995


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oxygen evolution is around 130 mV. This means that ϵ > 0 (eq 28), so that an increase in current density yields an increase in current efficiency for chlorine evolution (eq 26). That this is the case for chlorine evolution on DSA has been shown many times (refs 54, 56, 57, 62, 66, 69, 70, 75, 131, 154, and 162−164). Tilak et al. then went on to perform experiments to test these equations. The experiments were performed at T ≈ 95 °C using an electrolyte of 5 M NaCl and 2 M NaClO4 saturated with chlorine at pH ≈ 4. Coatings with decreased j0,oxygen evolution (RuO2 coatings doped with different amounts of HfO2 or SnO2) were tested using polarization measurements, showing that the overpotential for chlorine evolution was usually higher on such electrodes than on standard DSA. On the basis of galvanostatic selectivity measurements at 2.32 kA/m2, the selectivity for chlorine evolution was decreased for most of the Hf- or Sncontaining coatings. It was also found that the current efficiency for chlorine evolution increased with increasing current density. The percentage of oxygen in the off-gas using anode materials with increasing roughness factor (increasing Areal/Ageometric) was then measured at the same current density. The results are shown in Figure 12. The roughness factor was changed by using

Figure 11. Dependence of the oxygen content (vol %) of the cell gas on RTO anode composition obtained by Buné et al.68 (plotted using data from the article68). The NaCl concentration was 4.27 M, pH ≈ 1.6, and the current density was 1.2 kA/m2. Adapted with permission from ref 68. Copyright 1986 Springer Science+Business Media.

In 1987, Spasojević et al.161 presented the optimization of a RuO2−TiO2 anode for use in chloride electrolysis. Optimal Ru− Ti ratios, loadings, and calcination temperatures to achieve an active, selective, and stable anode were presented. Of interest to the question of selectivity is that results (very similar to those of Buné et al.68) indicating a clear decrease in partial current density for OER for coatings with Ru content below 40% were given. The lifetime of the anode during electrolysis of a dilute chloride solution was also maximized at around 40% Ru. Two years later, Tilak et al.69 published an article examining the selectivity in the chlorate process. Tilak et al.69 derived a mathematical expression, based on Tafel slopes, for the current efficiency of a main reaction when two competing reactions are occurring simultaneously. The equations are valid when both reactions are activation controlled. Tilak et al. believe that the relevant oxygen-evolving side reaction in the chlorate process is anodic oxidation of water (reaction 11). They disregard the reaction discussed by Kotowski and Busse,159 reaction 23, since they did not think it was “satisfactorily demonstrated”. The approach in developing the equations was similar to that of Buné et al.66 However, a more extensive set of equations was developed. The central one was 1 em = 1 + δem−ϵI −ϵ (26)

Figure 12. Relationship between surface roughness and the percentage of O2 in the gas phase, from the study of Tilak et al.69 Reproduced with permission from ref 69. Copyright 1988 The Electrochemical Society.

where I is the overall current density, em the current efficiency for the main reaction, j0, p δ= × e(k ′×ΔE) × j0,m A ·ϵ j0,m (27)

different coating deposition methods (brushing or spraying), different drying and calcination temperatures, and by varying the number of coating layers applied. The measurements showed that an increase in roughness factor gives an increase in the oxygen percentage in the off-gas. This seems to support the idea expressed by Uzbekov et al.142 and Trasatti and Lodi,165 that an increase in surface area yields a decrease in local current density and thereby a decrease in selectivity for chlorine evolution. In 1989, Hardee and Mitchell162 published an article regarding the selectivity for oxygen evolution at chlorate conditions, examining the effect of process conditions on the percentage of oxygen in the gas phase, using an undivided cell with a DSA. The electrolyte contained NaCl (either 50 or 200 g/dm3) and 2 g/ dm3 sodium dichromate at pH = 6.5 and 60 °C. Hardee and Mitchell found that hypochlorite decomposition was the main source of oxygen evolved. The percentage of oxygen in the evolved gas was dependent on the hypochlorite concentration at both sodium chloride concentrations. At the higher NaCl concentration, the dependence was approximately linear (see Figure 13). Extrapolation to cNaClO = 0 indicated that

and ϵ=

k − k′ k

(28)

k and k′ are the inverse of the Tafel slopes for the main reaction (chlorine evolution) and the “parasitic” reaction (oxygen evolution), respectively. ΔE is the difference in reversible potential between the main reaction and the parasitic reaction. A is the effective surface area of the electrode. The main implication of the expressions is that the selectivity for the main reaction (chlorine evolution in the case of chlorate and chloralkali production) could be increased if materials with a lower j0,p (j0,oxygen evolution) were used (eq 27). Furthermore, for chlor-alkali and chlorate processes, k > k′ since the Tafel slope for chlorine evolution on DSA is around 40 mV while the Tafel slope for 2996


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evolution rate at increasing concentrations of dichromate. Looking at the data presented by Hardee and Mitchell, a slight increase in the percentage of oxygen with increasing concentration of dichromate can be seen. Finally, long-term data for two anodes were reported: one of standard DSA and one with a secret modified composition. The modified coating, which had been adjusted based on the importance of the hypochlorite oxidation at the anode, showed a significantly lower percentage of oxygen in the off-gas. In the same year, Zhinkin et al.163 contributed a study mainly focused on stability of DSA. The work will be discussed in more detail in section 2.3, but it will be noted here that the authors showed that during chlor-alkali electrolysis using the conventional DSA and concentrated NaCl solution, the amount of oxygen in the off-gas increased with pH in the range of 2−5 (possibly related to the pH dependence of the O2/H2O equilibrium) and decreased with increasing current density in the range from 2 to 10 kA/m2. 2.2.4. Research during the 1990s. In 1990, Boodts and Trasatti166 published an article investigating the activity of Ru0. 3Ti(0.7−x)SnxO2 coatings for oxygen evolution from 1 M HClO4. The selectivity was not investigated per se but rather the activity and surface charge of this ternary oxide. Boodts and Trasatti found that the activity, measured as j(1.43 V vs RHE)/q* (current normalized by the surface charge), for oxygen evolution is increased by increasing Sn addition. In contrast to the suggestion of Krstajic et al.,73 Boodts and Trasatti believed addition of Sn to the coating should decrease the selectivity for chlorine evolution, and they proposed that it was PdSn2 rather than Sn itself that promoted the selectivity for chlorine evolution seen in the study of Krstajic et al. Also in 1990, Couper et al.167 reported further data for an ICI membrane chlor-alkali process anode coating. The coating, with unnamed composition, was an improvement of the previous low oxygen coatings154,158 and had been developed for a high oxygen evolution overpotential. Experimental data comparing the new low oxygen anodes with previous coatings indicated significantly lower oxygen percentages in the off-gas when using the new anode coating, at least up to pH ≈ 4.5. The results were explained using eq 29,

Figure 13. Dependence of hypochlorite concentration on the percent oxygen evolved as found by Hardee and Mitchell.162 The concentration of sodium dichromate was 2 g/dm3, the temperature was 60 °C, and the current density was 2.5 kA/m2. DSA were used. Reproduced with permission from ref 162. Copyright 1989 The Electrochemical Society.

the background oxygen production level, due to water oxidation, would be equivalent to 0.3% oxygen in the gas phase, for the 200 g/dm3 NaCl electrolyte. To determine the rate of homogeneous oxygen evolution, the cell was also operated in the same way but without the electrodes inside. A nitrogen flow, of similar magnitude to the hydrogen evolved in normal operation, was led through the cell to remove the formed oxygen. It was seen that the oxygen evolution was dependent on the NaClO concentration also in this case. The sum of the oxygen percentages from water oxidation and homogeneous hypochlorite oxidation was not enough to cover the total percentage of oxygen in the gas phase. Thus, the remaining part of the oxygen evolved was thought to be produced in reaction 23, as proposed by Kotowski and Busse.159 Hardee and Mitchell162 also examined the effect of some other process parameters. Increasing NaCl concentrations, from 50 to 300 g/dm3, were seen to reduce the oxygen evolution rate. The oxygen percentage in the off-gas decreased linearly with increasing NaClO3 concentrations between 0 and 400 g/dm3. It was thought that this might be due to competitive adsorption of chlorate onto the anode. There was no consideration of the explanation of Hammar and Wranglén,54 who showed that the increased viscosity of the electrolyte with increasing chlorate concentrations was a likely reason for the decreased oxygen evolution rate. Higher temperatures were seen to yield higher rates of oxygen evolution. This was partly explained by improved kinetics for homogeneous hypochlorite oxidation and partly by hypochlorite oxidation having a higher increase in rate with temperature than chloride oxidation. pH apparently had no effect in the 5−7 range, but an increased oxygen evolution rate was noticed at higher pH. This was thought to be related to the increasing prevalence of the deprotonated hypochlorite species at higher pH. The selectivity for oxygen evolution was seen to decrease with increasing current density. This was thought to be related either to difference in Tafel slope between the chlorineand the oxygen-evolving reactions or due to a limiting current density for hypochlorite oxidation being reached. Hardee and Mitchell also examined the effect of dichromate concentrations between 0 and 5 g/dm3. At 0 g/dm3, the oxygen percentage increased significantly. However, this was due to the large decrease in cathodic current efficiency that is associated with complete removal of dichromate from the electrolyte. When this effect was corrected for, the dichromate amount was said to have no effect on the oxygen evolution rate. This contrasts with the results of the work of Jaksić et al.,137 who saw increase in oxygen

log(oxygen content) = −Aη(O2 ) + B pH

(29)

which describes the oxygen evolution in membrane cells where the oxygen content of the cell gas is related partially to the pH and partially to the overpotential for oxygen evolution. A and B are introduced simply as fitting parameters without any physical interpretation. In the same year, Bergner et al.41 reported data for HCl addition to membrane chlor-alkali cells. This addition is done to neutralize any OH− that might migrate from the cathode side of the membrane. Bergner attempted to obtain data on suitable acid additions to the brine inlet of an industrial membrane chlor-alkali cell. The study did not attempt to decide the magnitude of different oxygen-producing reactions but presented suggestions for HCl addition that were shown to reduce the oxygen levels significantly, from just below 2% to 0.4% (see Figure 14). These results agreed with those previously obtained by Gorodetskii et al.151 Also in 1990, Wanngård57 reported results concerning the oxygen forming reactions in chlorate electrolysis in a pilot scale system, with industrially relevant current densities and temperatures and using a RTO anode and a steel cathode. Wanngård found, in agreement with other studies,54,162 that the oxygen 2997


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of the oxygen evolved originated from homogeneous hypochlorite decomposition (reaction 14). In 1992, Wanngård168 published a report concerning impurity effects in the chlorate process, where it was noted that electrolyte concentrations of Pb, Mn, Sn, F, Al, Si, and BaSO4 should be minimized to prevent deactivation (due, for example, to blockage) of the anode. Tests of hypochlorite decomposition catalysts under chlorate conditions at 70 °C had also been carried out, finding that Ni, Cu, Co, and Ir increased the oxygen percentage in the cell gas upon addition, while Pt, Ru, and Fe had no effects on the oxygen percentage (the form of the metal, i.e., salt, metallic, or oxidic, was not indicated). Wanngård also pointed out some important areas for further studies of the selectivity, such as studying the amounts of oxygen forming at the anode versus in the bulk electrolyte, deciding criteria for sufficient electrolyte purity and regarding demands on analysis methods used to detect impurities, and finally further studies of which current concentration gives the optimum selectivity. Also in 1992, Traini and Meneghini169 presented experimental data for the percentage of O2 in a chlor-alkali membrane cell offgas for a De Nora “Low-oxygen coating”, which decreased the percentage of O2 by about 0.4% at pH 4.0 in comparison with that of a conventional DSA coating. No information was given about the composition of the low-oxygen coating, although it was pointed out that an improved distribution of chlorides at the surface of the anode could be used to prevent local low brine concentrations leading increased rates of oxygen evolution. In the same year, Czarnetzki and Janssen170 presented a study of hypochlorite production by galvanostatic electrolysis (j = 2.76 kAm−2 or j = 5.54 kAm−2) of dilute NaCl solutions (6 to 90 g/ dm3) in a lab-scale membrane cell at alkaline pH (in the range 8− 10), using a commercial RTO anode. The study focused specifically on the initial behavior of the cell, with measurements covering about 1 h of electrolysis. It was found that increasing the hypochlorite concentration resulted in a linear increase in oxygen formation rate, although an extrapolation of the trend to zero hypochlorite concentration resulted in a significant background oxygen formation rate. Increasing the initial NaCl concentration instead resulted in decreased oxygen formation rates. Furthermore, increasing the current density resulted in a slight increase in oxygen formation rate. It was concluded that the rate of chloride oxidation was mass transport limited in the study, and therefore, the effect of adding nonreactive anions was attributed to their effect as supporting electrolytes. The addition of supporting electrolytes reduced the migration mass transfer of chloride to the electrode surface, thus further reduced the limiting current density for chloride oxidation. However, it was also noted that the addition of ClO−3 , ClO−4 , or SO2− 4 seemed to catalyze the oxygen formation, as a concomitant decrease in anode potential was measured. The authors also estimated that the pH at the anode during electrolysis was about 2.0, based on the formation rate of oxygen and an assumed concomitant production of H+ ions. Based on this, they concluded that the electrochemical oxidation of ClO− should be negligible and suggested that HOCl should be the dominating species at this pH. However, according to Figure 1, it is more probable that at a pH of 2, HOCl will be converted to Cl2 and water. Finally, the authors suggested that their chlorate formation results could be explained if the formation of chlorate occurred through both the conventional chemical conversion of hypochlorite, as well as through a direct electrochemical oxidation of Cl− to ClO−3 . In 1993, Bergner and Hartmann36 published an article concerning oxygen evolution in membrane chlor-alkali cells.

Figure 14. Effect of anolyte acidification with HCl on current efficiencies in chlor-alkali production, from Bergner.41 Original caption reads “Current efficiencies of products and by-products in brine pH range from −0.8 to 2.0. The points (×) at pH 2 correspond to the means of Figure 5.”. Reprinted with permission from ref 41. Copyright 1990 Springer Science and Business Media.

concentration in the off-gas was related to the concentration of hypochlorite species (the sum of [HOCl] and [ClO−]). In turn, the concentration of hypochlorite species was found to be dependent on pH, temperature, ionic strength, and the cubic root of the current concentration (total current divided by the total electrolyte volume in the system). Furthermore, Wanngård found that the oxygen percentage was independent of the flow velocity in the electrode gap. This contrasts with results obtained in studies at low temperature,53,54 where the oxygen evolution rate was mass transport limited and indicates that the oxygen evolution was kinetically limited at the higher temperatures used industrially. Wanngård called %O2,off − gas [HOCl + OCl−]electrodegap

= SR (30)

a “selectivity ratio” (SR) when discussing the results. Linear relationships (indicating that the rate of oxygen evolution is linearly proportional to the amount of hypochlorous acid species in the cell) were found when plotting SR versus (1/t), where t is the temperature in degrees centigrade. The same slope was found in all cases, except for when the electrolyte was contaminated by compounds active for hypochlorite decomposition. Wanngård claimed that this linear dependence implied reaction control of the oxygen evolution under chlorate conditions. Neither changing the electrode gap from 2 to 3 mm nor increasing the height of the electrode blades (i.e., changing from rectangular blades with 0.333 m height and 0.379 m width to rectangular blades with 0.8 m height and 0.3 m width) altered the slope of the linear dependence. However, the increase in height decreased the value for SR, implying that the electrode height (or possibly the total area, as the higher blades had a higher geometric surface area) had an impact on the oxygen formation. Wanngård cautioned that the effect on SR that higher electrodes exhibited could point to a difficulty in drawing conclusions relevant for full scale cells from laboratory scale data. Wanngård further claimed that experimental studies performed at the same company, at similar temperatures (50 to 90 °C), had shown that only 5−10% 2998


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in this review. Furthermore, Trasatti reviewed some data that points to the importance of a connection between surface redox reactions of electrocatalyst materials (e.g., changes the oxidation state of Ru in the surface coating of DSA) and the redox potentials of reactants. Trasatti comments that the reason why Tilak et al.69 and Spasojević et al.71 claim that Sn addition increases the selectivity for chlorine evolution is that “...the redox transition leading to oxygen evolution on RuO2 is pushed to more positive potentials by the presence of the inactive component. This is also the case for RuO2+TiO2 electrodes”. In 1994, further lab-scale research concerning the current efficiency in the chlorate process was presented by Baolian and Wenhua.173 The main contribution of the study was a new formula for calculating the current efficiency for a chlorate cell utilizing an oxygen cathode instead of the conventional hydrogen cathode. The study again reconfirmed previously seen trends regarding optimal pH for chlorate production (finding an optimum at around pH 6.5−6.75), sodium chloride concentration (increased efficiency with increasing NaCl concentration), and sodium chlorate concentration (a slight increase in current efficiency with increasing NaClO3 concentration). However, the study also examined the effect of additions of Na2HPO4, Na2Cr2O7, or NaF to the electrolyte at two different concentrations of NaCl (either 200 g/dm3 or 50 g/dm3). Additions of increasing amounts, up to 7 g/dm3, of Na2HPO4 or Na2Cr2O7 resulted in a slight decrease in current efficiency for chlorate formation. However, the addition of NaF resulted in increases in current efficiency of about 1% at the high NaCl concentration and about 3% at the low NaCl concentration. These results agree with the general conclusion of Fukuda et al.,144 who found that F− addition could reduce the OER selectivity of DSAs, while they disagree with those of Jaksić et al.,137 who found no effects on current efficiency for chlorate formation from fluoride addition. In 1997, Eberil et al.164 examined the selectivity issue in a chlorate system, calculating the efficiency for chlorate electrolysis based on the equation

Results were presented for oxygen and chlorate production in industrial membrane cells, operated at industrially relevant process parameters over a period of several years. The cells were either equipped with “low oxygen” or “high oxygen” anodes, with unnamed compositions. The results confirmed the previous trends141,150,154 in that the anode designed for lower oxygen production exhibited a lower oxygen production but also a higher chlorate production. In the same year, the first research, to our knowledge, that employs mass spectrometry (MS) for detection of Cl2 and O2 formed during electrolysis using DSA was published. Takasu et al.171 adapted DEMS for use with a titanium net coated with different compositions of RuO2, TiO2, and IrO2. The potentials where chlorine and oxygen, respectively, were first detected by the MS were termed the threshold electrode potentials. The difference between the two threshold potentials was proposed as a measure of the selectivity for the two reactions (see Figure 15).

Figure 15. Threshold electrode potentials of (a) oxygen and (b) chlorine evolution on RuO2−TiO2 electrodes of varying composition. Reprinted with permission from ref 171. Copyright 1993 Elsevier.

Also, the ratio between chlorine evolved and the total gas evolved at a certain potential (above the threshold potential for both reactions) was proposed as a measure of selectivity. Based on the results, the authors conclude that a 20% RuO2 + 80% TiO2 coating shows the highest selectivity for chlorine evolution. The applicability of the results to actual anodes in industrial production is uncertain, since DEMS studies are always done at low current densities. In this case, the maximum current density reached during the sweep was about 40 A/m2. In the same year, Nakajima et al.172 presented a study on the properties of Ru−Ti−Sn oxide and Ru−Sn oxide electrodes used for chlorine evolution in a laboratory-scale membrane cell. XRD indicated that the Ru−Sn oxide was of the rutile type, but the structure of the other electrodes was not clearly stated. The Ru− Sn oxide electrodes exhibited lower anode potentials than both Ru−Ir−Ti, Ru−Ti, and Ru−Ti−Sn electrodes. Similarly to conventional RTO electrodes, the volume percent O2 in the offgas increased with the amount of Ru in the Ru−Sn oxide coating. Lower Ru percentages in the mixed Ru−Sn coating also led to longer coating lifetimes, which also matches the behavior of RTO. In 1994, Trasatti’s latest extensive review of DSA and related conductive metal oxides was published.26 While the review discussed both chlorine and oxygen evolution on DSA and other materials, there is very little said about the selectivity between the two reactions. The main reference is to the studies of Krstajic et al.73 and Spasojević et al.71 on a selective water electrolysis and chlorate production anode, which has been described previously

ΦeNaClO3 = 100 − 2V (O2 ) − V (Cl 2) − 2.5

(31)

where 2.5 was a previously determined efficiency for cathodic reduction in the cell and V are volume fractions. Experiments under chlorate conditions showed that ηNaClO3 increased with potential, reaching a plateau at anode potentials E > 1.37 V versus SHE. The efficiency was constant at pH 6−7 at potentials above 1.365 V versus SHE. At potentials lower than the plateau value, the efficiency was significantly higher for solutions with pH 6− 6.5. The maximum ΦeNaClO3 was found at potentials close to the so-called “critical potential”, Ecr. This potential has been associated with a break in the Tafel line and an increase in the rate of decomposition of the anode.165,174 Furthermore, they found that the efficiency of chlorate production decreased with increasing temperature but increased with increasing current density. Thus, a high current efficiency can be maintained at a higher temperature if the current density is increased. In 1998, Arikawa et al.75 published DEMS results on the selectivity between oxygen and chlorine evolution on RTO and RuO2. Several known effects, regarding NaCl concentration, oxide loading, and RTO composition, were reconfirmed by their experiments. The experiments using different concentrations of NaCl showed that the selectivity for chlorine evolution increased with increasing concentration of NaCl. Additionally, the selectivity increased with increasing anode potential (and current 2999


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increasing temperature. At Ea > Ecr, ΦeO2 increased with increasing temperature. The efficiency for chlorate production was determined largely by ΦeO2. A higher temperature, a lower pH e and a higher concentration of NaCl increased ΦCl . The 2 measured efficiencies were also compared to those obtained in the system164 with directed circulation and a holding tank for chlorate formation. The system with the holding tank had higher efficiencies and showed similar trends in efficiency at E < Ecr. No experiments at E > Ecr had been done in the system with the holding tank and directed circulation. Furthermore, a difference between the Ir-containing DSA and the standard DSA was noted. For the standard DSA, the maximum ΦeNaClO3 was moved to higher anode potentials (higher applied current densities), and ΦeO2 increased overall, when the chloride concentration was decreased. This could be interpreted as resulting from concentration limitations of Cl− at the anode surface. However, when Ir-containing DSA were used, ΦeO2 did not increase as the chloride concentration was decreased. The reason for the increase in oxygen evolution at standard ruthenium−titanium DSA, once the critical potential is reached, was supposed to be the formation of volatile RuO4. As IrO2 increases the stability of the coating, it was proposed that this process is hindered on IrDSA. The Ir-DSA coating showed other advantageous characteristics, such as somewhat lower ΦeCl2 and higher ΦeNaClO3, at most current densities. In 2001, Byrne et al.176 published a tertiary model of current distribution in the chlorate cell, including mass transport through diffusion, migration, and convection for relevant species, using the Nernst−Planck equation. Butler−Volmer expressions were used to model reactions 2, 11, and 23 on the anode. The last reaction was modeled for oxygen formation through both electrochemical oxidation of hypochlorite and electrochemical oxidation of hypochlorous acid. The possibility of different type of heterogeneously catalyzed, nonelectrochemical, decomposition of hypochlorite on the surface of the electrode was not considered. Byrne found that electrochemical water oxidation was an insignificant source of oxygen when compared with that of anodic decomposition of hypochlorous acid or hypochlorite. The latter reactions were the dominating source of oxygen at the anode. Two different trends were seen for the anodic decomposition of hypochlorous acid and the hypochlorite ion. While the decomposition of hypochlorous acid was highly dependent on the flow rate, the rate of decomposition of hypochlorite ions was found to be relatively independent of the flow rate. This agreed well with the results of Wanngård,57 indicating that hypochlorite, rather than hypochlorous acid, might be the main reactant in the anodic decomposition of hypochlorite species in the chlorate process. The model was experimentally validated using a cell with segmented electrodes in a vertical column.177 It was found that the model corresponded well to the experimental data only at low flow velocities between 0.25 and 0.75 m/s and near the inlet inside the column. The same year, Jirkovský et al.76 published a study concerning sol−gel-prepared RuO2 coatings on Ti meshes. The study was focused on the selectivity for chlorine and oxygen evolution on coatings made up by crystals of different sizes. The particle sizes increased with the annealing temperature used during the preparation. The DEMS technique was used to measure the oxygen and chlorine evolution during cyclic voltammetry (CV) measurements. The study focused on examining the activity of different crystal edges on RuO2 nanocrystals. The electrolyte was

density) in solutions with high chloride concentrations (2 M). A number of measurements were also carried out at a lower chloride concentration of 0.3 M. At this concentration, the selectivity for ClER in general decreased with increasing anode potential. The decrease was more severe when the oxide loading of RuO2 was increased from 0.1 to 0.9 mg/cm2. For a fixed oxide loading of RTO, it was found that the selectivity for ClER was improved over that of pure RuO2 when the RuO2 content in the coating was between 20% (the lowest concentration) and 60%. 2.2.5. Research during the 2000s. This section is divided into one part describing experimental studies and one part devoted to a more detailed description of the first-principles modeling of chlorine and oxygen evolution on metallic oxides. This latter field has seen a large growth during the past few years. In the present review, these studies are described after the experimental studies, although the approach used in some of the experimental studies has been inspired by recent results from first-principles calculations. 2.2.5.1. Experimental Studies. In 2000, Eberil et al.74 published a thorough examination of the current efficiency for chlorate production, by measuring the evolution of chlorine and oxygen, using both DSA (30 mol % RuO2 + 70 mol % TiO2) and Ir-containing DSA (15 mol % RuO2 + 15 mol % IrO2 + 70 mol % TiO2). The maximal ΦeNaClO3 was again found at potentials E ≈ 1.4 V versus AgCl/Ag, close to Ecr (see Figure 16). This coincided with the minimum in ΦeO2. Depending on whether the anode potential was below or above the critical one, the dependence of the efficiency for oxygen evolution, ΦeO2, on the anode potential, Ea, was different. At Ea < Ecr, ΦeO2 decreased with

Figure 16. Current efficiencies for O2 and Cl2 formation as a function of anode potential, determined under different conditions. Unfilled, halffilled, and filled symbols depict measurements at 60, 80, and 90 °C, respectively. The series of symbols (circles, squares, triangles, and stars with increasing number of points) depict measurements done at different current densities from 0.15 to 1.7 A/cm2. Original caption reads (Eberil et al.74 used the symbol η to signify current efficiency, rather than using the conventional Φe175): “Dependences of ηCl2 and ηO2 on Ea of ORTA at different ia and T in solution containing 150 g l−1 NaCl and 350 g l−1 NaClO3; pH 6.5.” ORTA is an abbreviation meaning ruthenium−titanium oxide anode (RTO). Reprinted with permission from ref 74. Copyright 2000 Springer Science and Business Media. 3000


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thought to be active for oxygen evolution, sites with one nickel atom were thought to be active for chlorine evolution, and sites with two Ni atoms were imagined to be selective for oxygen evolution. The third type was thought to be less active for oxygen evolution than the first type. This model could describe some trends observed in the experimental data but was unable to account for the low activity for chlorine evolution that coatings with the highest Ni content exhibited. There was no comment on the stability of the coating, but Pourbaix diagrams45−47 indicate that Ni2+ might dissolve from the coating. Also in 2008, Macounová et al.180 also presented results for Fedoped RuO2 (Ru0.98Fe0.02Ox and Ru0.7Fe0.3Ox) prepared using sol−gel synthesis. XRD indicated that the Fe dopants entered into the overall Ru rutile lattice. Selectivity measurements using DEMS indicated that the selectivity was affected strongly by the gas present in the 0.15 M NaCl−0.1 M HClO4 electrolyte. In the Ar-saturated electrolyte, the selectivity for Cl2 was high for both electrode compositions, whereas in oxygen-saturated electrolyte, oxygen evolution was the dominating process for the Ru0.7Fe0.3Ox electrode. The authors suggested that the trend could be explained by oxidation of the Fe sites in oxygenated solutions. The same group also published an article concerning the Ru0.8Co0.2O2−x material in 2008.78 The electrodes were prepared by sol−gel synthesis onto Ti meshes. The study used X-ray photoelectron spectroscopy (XPS) and TEM to draw conclusions regarding the connection between the surface morphology and the selectivity between chlorine and oxygen evolution on this material. DEMS measurements were done in the same way as the previous articles from the same group, by using CV and DEMS. This time, the solution not only contained 0.1 M HClO4, but also (like in refs 76 and 77) 0.3 M NaCl. The addition of chlorides shifted the onset potential for oxygen evolution to more anodic potentials. The particle size did not have any effect on the rate of oxygen evolution, but an increased selectivity for chlorine evolution (due to an increased rate of chlorine evolution) was noted on coatings with larger (d > 40 nm) crystals. The onset potential for the chlorine evolution reaction was about 1.1 V versus SCE, while the onset potential for OER was approximately 1.2 V. At higher potentials, the chlorine evolution current was almost constant, while the oxygen evolution current continued to increase. The Tafel slope for the chlorine evolution was 120 mV/dec below 1.2 V and a very high value of about 600 mV/dec at higher potentials. The change in Tafel slope at 1.2 V was thought to be related to a change in the number of active sites for chlorine evolution at that potential. Makarová et al.78 claimed that this supported the Krishtalik mechanism, in which an oxidation of the surface is the first step in the chlorine evolution reaction. The oxygen evolution reaction was thought to be retarded on the mixed Ru−Co oxide, due to stabilization of the surface oxygen intermediate by Co. Two papers concerning ZnII-substituted RuO2 were published in 2010 and 2011.181,182 The coating was prepared by use of freeze-drying, with subsequent high temperature annealing and deposition onto Ti meshes. The coating was characterized by use of DEMS (in 0.1 M HClO4 and 0.15 M NaCl), X-ray absorption spectroscopy (XAS), X-ray absorption near-edge spectroscopy (XANES), extended X-ray absorption fine-structure spectroscopy (EXAFS), EDX, scanning electron microscopy (SEM), and transmission electron microscopy (TEM). XAS and EXAFS data showed that addition of Zn into the structure obstructed the stacking of Ru coordinatively unsaturated sites (CUSs). On the basis of refined EXAFS data, it was proposed that Zn gathered in

a dilute solution of 0.1 M HClO4 and 0.3 M NaCl. For the coatings with the smallest particles, Jirkovský et al. found that the onset potential for chlorine evolution was more positive than that for oxygen evolution, as opposed to what was found in the study of Takasu et al.171 However, Arikawa et al.75 also made measurements at cNaCl < 0.5 M, finding that the threshold potential (or onset potential) of oxygen evolution was lower than that of chlorine evolution. Furthermore, Jirkovský et al.76 found that the Tafel slopes of both oxygen and chlorine evolution were 114 mV/dec. These values are, like in the previous studies using DEMS,75,171 much higher than what is normally reported for RuO2. For coatings prepared at high temperature, with larger particles, the threshold potential for oxygen evolution shifted to a higher potential than that for chlorine evolution. The selectivity for chlorine evolution was higher on coatings prepared at these higher temperatures and showed a linear dependence on the annealing temperature. The authors rationalized the increased selectivity for chlorine evolution based on the types of crystal edges that were more prevalent at the larger particles that formed at these temperatures. However, the specific surface area of the coating was about half as large on the electrodes prepared at 900 °C. It is known from previous studies that an increase in current density increases the selectivity for chlorine evolution. Whether this had any effect does not appear to be considered in the article. Also in 2006, Santana and De Faria178 presented research on the electrolytic production of Cl2 and O2 using DSA doped with CeO2 and Nb2O5. Their conclusions were that Nb2O5 had no effect on the activity or selectivity of the coatings. However, introduction of even 1 mol % CeO2 into DSA decreased the low Tafel slope for O2 evolution to 30 mV/dec from 40 mV/dec, but only at current densities lower than 10 A/m2. On the other hand, CeO2-doping did not seem to affect the Tafel slope for Cl2 evolution. CeO2-doped DSA thus exhibited higher O2-selectivity at low current densities. In 2007, Panić et al.179 presented stability and selectivity data on RTO deposited, by use of a sol−gel procedure, on a ternary titanium carbide, Ti3SiC2. It was thought that this carbide would be much more stable as a support material than Ti. Instead of the more usual 30% RuO2 + 70% TiO2, a coating with equimolar amounts of ruthenium oxide and titanium oxide was used. Experimental results suggested that the selectivity for chlorine evolution was more favorable for anodes on the Ti3SiC2 support than on Ti. However, the new support was shown to be much less stable in an accelerated lifetime test than the standard Ti support. In 2008, Macounová et al.77 published DEMS results on the selectivity between further DEMS research examining the selectivity between oxygen and chlorine evolution on Ru1−xNixO2−y, where x was between 0.02−0.3, coated on Ti meshes. Similarly to the study by Jirkovský et al.,76 the coatings were characterized using XRD and energy-dispersive X-ray spectroscopy (EDX) and evaluated using DEMS and CV in solutions containing 0.1 M HClO4 and 0.3 M NaCl. This time, the coatings were all heat treated at 400 °C for 20 min. The DEMS study indicated a complicated dependence of the selectivity on the Ni content of the coating. The selectivity was dependent both on the applied potential and the Ni content. The maximum selectivity for chlorine evolution was found at x = 0.1. The Tafel slope for chlorine evolution was found to be around 35 to 40 mV/dec, while that for oxygen evolution assumed its lowest value of 40 mV/dec at x = 0.02 and then had a maximum of 100 mV/dec at x = 0.1. A model was presented to explain the activity, based on there being four types of active sites, with increasing numbers of adjacent Ni atoms. Sites with no Ni atoms were 3001


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solutions could then be explained by concurrent oxygen formation via both the peroxo route suggested by Rossmeisl et al.184 and by recombination of oxygen intermediates on crystal edges in the ilmenite structures on the surface of the coating. It was proposed that at x < 0.1, the recombination mechanism could make up for the lost rutile structure peroxo species formation. In two articles published in 2011 and 2012, Chen and coworkers studied the seawater electrolysis performance of RuO2− Sb2O5−SnO2 electrodes with and without IrO2 in the coatings.185,186 The Ir-doped coatings were found to be smoother, and their stability in an accelerated life test was found to be higher. Current efficiencies were similar for both electrodes. In 2012, Neodo et al.187 compared the electrolytic behavior of RuO2−SnO2-coated Ti electrodes with platinized Ti electrodes. The study was performed in dilute NaCl solutions (0.07 M). Current efficiencies for active chlorine, O2, chlorate, and perchlorate were determined. It was found that the current efficiency for oxygen formation increased with increasing temperature (tests were performed at 10, 25, and 65 °C), consistent with a change in activation energy for oxygen-evolving reactions. Furthermore, the current efficiency for oxygen formation was slightly higher for the platinized Ti electrode. In the same year, Spasojević et al.60 presented current efficiency and stability studies for a RTO electrode with an overlayer of Pt and rutile IrO2 nanoparticles. The electrode was prepared on Ti, by first thermally depositing a 10 g/m2 40% RuO2-60% TiO2 layer, and then thermally depositing a second layer with a loading of 6 g/m2 from 60% H2PtCl6 and 40% IrCl3 solution. This electrode was compared with an electrode with only the first RTO layer. The double-layer electrode was found to exhibit 50 mV higher overpotentials for oxidation of dilute pH 7 NaCl solutions than the more standard RTO electrode. Furthermore, the double-layer electrode had a much smoother surface and a significantly lower electrochemically active surface area. The anodic current efficiency for active chlorine formation (evaluated based on the gas composition) was some 4% higher for the double-layer coating than for the single-layer RTO coating (see Figure 18). By extrapolating to zero hypochlorite concentration, it could be concluded that OER is the dominating source for oxygen evolution for both coatings. Furthermore, the stability, as evaluated by the time before a steep increase in electrode potential after galvanostatic electrolysis at j = 10 kAm−2 of a 0.5 M H2SO4 solution, was found to be significantly higher for the double-layer coating than for the single-layer RTO coating and several times higher than for a DSA electrode of conventional composition. In 2013, Xiong et al.188 presented research on DSA coatings and Sn−Sb-substituted DSA coatings deposited either on Ti or on TiO2 nanotubes (TNT). In the substituted DSA coating, half of the Ti content had been replaced by Sn and Sb to obtain coatings of the nominal structure Ru 0.3Ti0.34Sn0.3Sb0.06O2 (Ru0.33Ti0.4Sn0.18Sb0.08O2 according to EDX). SEM imaging showed that the coatings deposited on TNT were smooth, in contrast to the usual mud crack structure of coatings on Ti. Determination of the voltammetric surface charge for the coatings, including one coating of overall composition Ru0.3Ti0.34Sn0.36O2, showed that the Sn−Sb-substituted coating deposited on TNT had the highest electrochemically active surface area, slightly higher than that of both RTO deposited on TNT and Ru0.3Ti0.34Sn0.36O2 deposited on TNT. It was suggested that the introduction of Sb possibly retards crystallite growth, yielding a higher electrochemically active surface area.

ilmenite defects on the surface, similarly to how Fe impurities in rutile TiO2 gathers. A graphical representation of rutile and ilmenite RuO2 and Ru0.9Zn0.1O2 from the article by Petrykin et al.182 can be seen in Figure 17. The EXAFS function for ilmenite

Figure 17. Graphical representations of rutile and ilmenite RuO2 and Ru1−xZnxO2 from Petrykin et al.182 It is seen that the ordered arrangement of Ru CUSs (indicated by adsorbed O in the figure) is obstructed in the ilmenite structure. Reprinted from ref 182. Copyright 2011 American Chemical Society.

structure Ru1−xZnxO2, based on the structure of ZnTiO3 ilmenite but with Ti atoms replaced with Ru atoms in the optimization of the structure, fit the measured Ru-EXAFS function well. Since no ilmenite structures were indicated in XRD or electron diffraction patterns, it was thought that they should be present in small domains distributed in the catalyst crystallites. No information was given regarding the stability of Zn in the coating, but Pourbaix diagrams45−47 indicate that ZnII might leach from RuO2 at low pHs. DEMS data showed that in solutions without chlorides, the activity of Ru0.9Zn0.1O2 was higher than that of RuO2, based on a lower onset potential and a higher current density at three potentials above the onset potential. However, at higher levels of Zn addition, the activity started to decrease while staying slightly higher than that of pure RuO2. In DEMS measurements in solutions containing chlorides, the activity for concurrent oxygen and chlorine evolution was lower than that of pure RuO2. While a similar maximum of activity for 10% substitution was observed also in this case, the current density was lower than that of RuO2 for all degrees of substitution and at all applied potentials. However, the ZnII-substituted coating had a much improved selectivity for oxygen evolution and the selectivity increased with increasing substitution by Zn. The increased selectivity was rationalized based on the presence of ilmenite structures in the surface of Ru1−xZnxO2. The formation of peroxo groups, found to precede the formation of Cl2 on RuO2 according to the work by Hansen et al.183 (see the following section), is not favorable since CUSs and bridge sites are distributed randomly on the ilmenite structures (see Figure 17). However, this blocks the formation of precursor intermediates for both chlorine and oxygen evolution. This would conflict with the observation that the activity for oxygen evolution was improved by substitution with Zn in chloride-free solutions. Therefore, Petrykin et al.182 thought that oxygen evolution could take place via crystal edge recombination of surface oxygen intermediates on the ilmenite structures. The higher activity of Ru0.9Zn0.1O2) in chloride free 3002


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reaction barrier for OER, offering a potential explanation for the trend in selectivity observed for Co- and Mn-doped RuO2. Later the same year, Krtil and co-workers also examined Co, Ni, and Zn-doped IrO2 in an analogous way.189 On the basis of XRD and EXAFS measurements, it was concluded that Zn and Co were doped into the cationic positions of the rutile lattice, while Ni formed clusters embedded in the overall structure. The concomitant evolution of oxygen and chlorine during electrolysis was characterized using DEMS in experiments where either the concentration of NaCl or the anode potential was varied. Increasing the chloride concentration yielded the expected increase in selectivity for chlorine evolution. The pure IrO2 material showed a higher selectivity for chlorine evolution than any of the three-doped materials. Whereas the doped materials showed decreasing Cl2 selectivity at potentials between 1.55 to 1.65 V versus RHE and the highest NaCl concentration of 0.05 M, the pure material showed 100% Cl2 selectivity up to the maximum 1.65 V versus RHE potential. At 0.01 M NaCl, however, all materials showed similar and decreasing Cl2 selectivities of around 10−20% at potentials above approximately 1.6 V versus RHE. At potentials between 1.4 to 1.6 V versus RHE, the Cl2 selectivity of the pure IrO2 material was significantly higher than that of the doped materials. In 2015, Le Luu et al.190 studied the selectivity between Cl2 and O2 evolution in acidic 5 M NaCl solution, comparing IrO2 and RuO2 electrodes. The electrochemically active surface areas of the electrodes, which were prepared by conventional thermal decomposition, were determined using CV. The IrO2 electrodes were found to exhibit a 40% larger active surface area. Electrochemical measurements at pH 2, 4, and 6 showed that the current efficiency for chlorine evolution was higher for the RuO2 electrodes. The current efficiency for both electrodes decreased as the NaCl concentration was decreased, but the RuO2 electrode in all cases exhibited a higher current efficiency. The authors suggested surface oxygen deficiency, XPS d-electron binding energies, differences in point-of-zero charge, electrical conductivity, and hydrophilic properties as reasons behind the increased selectivity of the RuO2 electrodes. However, it is perhaps more likely that the difference in selectivity was the result of the significantly lower electrochemically active surface area which the RuO2 electrodes were found to exhibit (see section 3.1.3 for further discussion). In 2014, Spasojevic et al.191 presented a new chlorate production current efficiency model. The chlorate cell was modeled as either an ideal tubular reactor or an ideal stirred tank reactor. The model only included losses due to water oxidation (reaction 11) and anodic chlorate formation (reaction 10). A DSA and a Ti plate cathode were used as electrodes. Experiments were performed in a pumped laboratory chlorate setup, using a conventional small-volume electrolysis cell connected to a larger holding tank, to get experimental data for the model. The conditions for the experiments were close to those used in industrial cells. The model was verified by comparing predicted and measured inlet and outlet active chlorine concentrations, for two different inlet concentrations and for electrodes of different sizes, yielding good agreement. This means that only the loss reactions of water oxidation and of anodic chlorate formation were enough to account for current efficiency losses in a laboratory chlorate production setup. In the same year, Zeradjanin et al.192 presented experimental results for the selectivity between chlorine and oxygen evolution on Ti−Ru−Ir mixed metal oxide electrodes. DEMS was used to determine the evolved gas composition. The DEMS measure-

Figure 18. Anodic current efficiency for active chlorine (hypochlorous acid and hypochlorite) production, based on measurements of O2 in the cell gas, as a function of hypochlorite concentration, from Spasojević et al.60 ■ are a conventional 40 mol % RuO2-60 mol % TiO2 coating, while ● represent the Pt-IrO2 double-layer coating. Experiments were carried out in a 30 g/dm3 NaCl solution at 25 °C and a current density of 1 kAm−2. Reprinted from ref 60. Copyright 2012 Elsevier.

Linear sweep voltammetry indicated that the onset potential difference between Cl2 and O2 evolution was highest for the Sn− Sb-substituted coating deposited on TNT, followed by the Sn− Sb-substituted coating deposited on Ti, indicating an enhanced Cl2/O2 selectivity. The increased Cl2 selectivity is thus a result of the introduction of the tin−antimony component into the coating, rather than an effect of the smoother electrode surface. Since previous studies have found Sn-doping to increase the Cl2 selectivity, it is possible that this component (and not the Sb component) is the main origin of the selectivity increase. The work of Krtil and co-workers continued in 2013 and 2014, now in cooperation with Rossmeisl and co-workers.82,83 The focus was now on RuO2 doped with small quantities (up to 20 mol %) of Co, Ni, or Zn. In the initial experimental study of Codoped RuO2, the structural and electrochemical properties of this mixed oxide were tested using XRD, XAS, EXAFS, and DEMS experiments examining both chlorine and oxygen evolution in 0.1 M HClO4 with 0.01 to 0.3 M NaCl. Refinement of the EXAFS functions indicated that the rutile structure was maintained, with Co occupying sites diagonally across the unit cell (i.e., with no neighboring Co CUS). The selectivity for oxygen evolution was found to increase with increasing concentration of Co in the coating. The relationship between the selectivity trend and the structure of the mixed oxide electrocatalyst was examined in the following study by Halck et al.83 In this study, electrochemical selectivity measurements on Co- and Ni-doped RuO2 electrocatalysts were combined with first-principles DFT calculations (see also the following section) based on EXAFS-refined structures. The calculations showed that the presence of Ni or Co dopants serve to activate the bridge site oxygen, allowing it to accept a hydrogen adsorbate. This modification of the OER mechanism, which on rutile materials involves the formation of an *OOH intermediate adsorbed on, for example, Ru alone, stabilizes the *OOH intermediate. The binding energy of the *OOH intermediate is determined by, for example, Ru alone on pure rutile materials, but in these doped materials, it can be tuned by both the main constituent cation and by the dopant cation. This stabilization of the *OOH species decreases the overall 3003


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ments were performed at 4 kA/m2, 25 °C, pH = 3 and NaCl concentrations varied between 1 and 4 M. Scanning electrochemical microscopy (SECM) was used to study the impact of the microstructure of the electrode on the activity and selectivity for the two reactions. DEMS indicated that oxygen was evolved even at NaCl concentrations of 3 M, over which the current efficiency for chlorine evolution reached almost 100%. EDX measurements found local Ru enrichment after electrochemical testing. SECM measurements, sensitive only to the evolved Cl2, found the current distribution over the electrode surface was inhomogeneous on the micrometer scale. The SECM tip current (which is due to the reduction of Cl2 formed at the electrode surface) varied from −500 to −3000 nA, and it was estimated that the local electrode potential at the most highly active site was 1.41 V versus AgCl/Ag and at the least active site 1.63 V versus AgCl/Ag. It was thus deemed possible that Cl2 could be formed at certain active parts of the electrode, only to be reduced at other, less active, parts. That this was possible despite the high anodic potentials was explained based on an increase in the Cl2/ Cl− redox potential due to the combination of a high concentration of formed Cl2 and a low concentration of Cl− (due to concentration limitations). In 2015, a purely experimental study on Mg-doped RuO2 was also published.193 From the group of Krtil, the work employed many of the same techniques as previously (XRD and EXAFS measurements for structural studies and DEMS for electrochemical studies). The authors concluded that Mg was unevenly distributed in the overall RuO2 rutile structure, and that the local environment surrounding the Mg dopants was ilmenite-like, similar to what was found for Zn-doped RuO2,182 except for the lowest dopant concentration. The Mg-doped material was found to exhibit improved Cl2 selectivity. This effect was most significant for the lowest dopant concentration of 5 mol %. The authors believed that the structure of this Ru0.95Mg0.05O2 electrode could be most similar to that of undoped RuO2, indicating that a significant improvement in selectivity could occur when Ru is replaced by Mg in the rutile lattice. Also higher concentrations of Mg dopant resulted in higher chlorine selectivity, but the rationale given was then the broken CUS stacking in the ilmenite structure.182 In the same year, Macounová et al.84 studied anodic oxidation of hypochlorite in 0.1 M NaClO4, at pH values of at least 9.5, using DEMS and a Pt mesh electrode. It was suggested that, at an anodically polarized Pt electrode, hypochlorite is decomposed in a radical reaction involving formation of ClO•. Once this species has formed, it was suggested that it should react with OH− to form the superoxide radical •OO−, which could either react with water to form H2O2 and O2 or with HOCl to form O2 and OH·. The latter radical was suggested to regenerate ClO−, closing a catalytic loop resulting in conversion of water into O2. On the basis of the electrical charge transferred per mol of O2 generated on the Pt electrode, which was lower than expected during OER, it was suggested that the reaction via ClO•, •OO−, and OH• was occurring on the electrode. This pathway does not involve formation of H2O2 and results in accumulation of Cl− in the system. However, H2O2 was detected, suggesting that also the reaction between •OO− and water was ongoing. Furthermore, measurements of chloride concentrations before and after an experiment did not indicate accumulation of Cl− and none of the suggested radical species were detected experimentally. Whether the suggested radical mechanism is relevant for industrial chloralkali or sodium chlorate production using DSAs remains an open question.

2.2.5.2. Research Based on First-Principles Modeling Concerning Chlorine and Oxygen Evolution on Metal Oxides. The number of articles using first-principles approaches to understand electrochemical processes has increased substantially since the turn of the century. The majority of studies performed so far fall into the group of computational free energy (or ab initio thermodynamics) studies,194 where only the overall thermodynamics of the elementary steps of surface reactions are characterized. The two articles which laid the groundwork for the first theoretical works on the selectivity between oxygen and chlorine evolution on metal oxides are those of Rossmeisl et al.184 and Hansen et al.,183 published in 2007 and 2010, respectively. The first article models oxygen evolution on a few metallic oxides, while the second one continues the work by also including chlorine evolution in the model. Both studies employed the “computational hydrogen electrode” (CHE) method of Nørskov et al.195 This method allows results from first-principles calculations in which only adsorption energies of intermediates on, for example, metal oxide slabs in vacuum are calculated to be used to describe electrochemical reactions in solution, by correcting the adsorption energies in vacuum using standard thermodynamic relationships and reference data. This significantly simplifies the computational expense and complexity of the modeling, as neither the charge transfer process nor the effect of polarization on the binding energies of intermediates have to be modeled explicitly. However, this also means that the full kinetics of electrochemical processes cannot be captured in the model. For this to be possible, methods including explicit polarization of the model surface, such as those of Filhol and Neurock,196 Otani and Sugino,197 or Rossmeisl et al.198 have to be used. Nevertheless, the existence of Brønsted−Evans− Polanyi (BEP) relationships in heterogeneous catalysis and electrocatalysis,199−201 which show that the activation energy for an elementary reaction step in general is correlated with the reaction free energy of the same step, does indicate that the conclusions from computational free-energy studies give a qualitatively correct description of the full thermodynamics and kinetics of surface reactions. Indeed, the successful predictions of more active electrocatalysts for a variety of reactions that have resulted from computational free-energy studies support this conclusion.202,203 The computational free-energy study of Rossmeisl et al.184 used DFT calculations to calculate the stability of intermediates in the oxygen evolution reaction on the (1 1 0) surface of RuO2, IrO2, and TiO2. The (110) surface of rutile oxides has two different sites on which reactions can take place, bridge sites or CUSs. The nature of the (1 1 0) surface of RuO2 was studied previously by Over et al.,204 and the importance of CUS for the activity of oxide materials was conjectured already by, among others, Mars and van Krevelen.205 A model of the (1 1 0) surface of RuO2 is found in Figure 19. Rossmeisl et al.184 calculated the adsorption free energies of all intermediates that could result in oxygen evolution. The calculations showed that oxygen evolution on rutile oxides occurs via the following steps, where HOc, Oc, and HOOc are adsorbates bound to the CUS:

3004

H 2O → HOc + H+ + e−

(32)

HOc → Oc + H+ + e−

(33)

Oc + H 2O → HOOc + H+ + e−

(34)

HOOc → O2 + H+ + e−

(35) DOI: 10.1021/acs.chemrev.5b00389 Chem. Rev. 2016, 116, 2982−3028

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Figure 19. A model of the (1 1 0) surface of RuO2. Ruthenium cations are green, while oxygen anions are red. The Ru CUSs and the O bridge sites (binding to Ru bridge sites) are indicated with arrows.

Rossmeisl et al. found oxygen evolution through recombination of two oxygen atoms (a Volmer−Tafelmechanism) to have a too high activation energy to be plausible, based on previous DFT calculations by the same group.199 This means that the traditional view of OER on rutile oxides as being dependent on O−O recombination on neighboring Ru CUS67,68 is likely incorrect. In the case of RuO2 and IrO2, the largest ΔG was found for the third step in the mechanism above, the decomposition of the second water molecule. Therefore, for RuO2 or IrO2, the calculated ΔG for this third step was used as the measure of the overpotential. The overpotential (ΔGstep 3 − ΔG0 = ΔGstep 2 −1.23 V) was determined to be 0.37 V for RuO2 and 0.56 V for IrO2. In the case of TiO2, the step associated with the highest change in free energy was the second one, formation of O* from HO* on the surface. The overpotential, based on ΔG for this step, was 1.19 V. The adsorption energies for the intermediates (Oc, HOc, and HOOc) for different levels of surface coverage of Oc and HOc on the surfaces of the oxides were also calculated. A linear relationship, a scaling relation, was found between the adsorption energy for O on the surface and the adsorption energy of HO or HOO. In other words, surfaces that bound O strongly also bound HO and HOO strongly. This also implied that, for all three oxide materials treated, the rate of the oxygen evolution reaction is given by the adsorption energy of oxygen on the surface. The scaling relation was then applied in the calculation of the free energy of reaction for the different steps of the reaction. This yielded a volcano plot, as shown in Figure 20. The volcano plot qualitatively agreed with experimental data (mainly that described by Trasatti27,156). The topic of the present review is the selectivity between chlorine (or chlorate) formation and oxygen evolution, and thus all studies on ways to improve the OER on metallic oxide materials will not be included here. However, it is interesting to compare the results of the study of Rossmeisl et al.184 with the later theoretical study of Fang and Liu.206 In 2010, Fang and Liu206 studied the OER on RuO2 (110), now also calculating the kinetic barriers for both the Volmer−Tafel pathway (recombination of two Oc) and a water dissociation pathway, similar to that suggested by Rossmeisl et al.184 The study represents a step forward as the calculation of kinetic barriers allows for an explicit account for both reaction thermodynamics and kinetics. Fang and Liu206 found, in agreement with the assumption of Rossmeisl et al.,184 that the Volmer−Tafel pathway was less energetically favorable than the water dissociation mechanism. A difference between the mechanisms of Rossmeisl et al.184 and Fang and Liu206 resided in whether the Obr sites were protonated. In both cases, a protonated bridge site was found to reduce the thermodynamic reaction barrier, but Rossmeisl et al.184 concluded that bridge sites will be deprotonated at anodic

Figure 20. Volcano plot for oxide materials obtained by Rossmeisl et al.184 The different lines represent the negative of the free energy of reaction of each step. These lines each limit the reaction rate, based on the binding energy of oxygen on the surface of the three oxides and the surface coverage. The most active type (based on surface coverage) of each oxide is included as ● in the diagram. Reprinted with permission from ref 184. Copyright 2007 Elsevier.

potentials above 1.4 V versus SHE. However, Fang and Liu206 found that a mixed HObr-Obr phase was slightly more stable than a fully deprotonated RuO2 surface at potentials up to about 1.6 V versus SHE, enabling the HObr-stabilized pathway. Interestingly, in 2014, a type of Obr stabilization was suggested by Halck et al.83 to explain the improved OER activity of Ni- and Co-doped RuO2. In that case, it was suggested that the dopant cations predominantly occupied the cationic bridge sites, allowing the Obr sites to stabilize the HOObr intermediate in the mechanism of Rossmeisl et al.184 In 2010, Hansen et al.183 applied DFT calculations to treat chlorine evolution at rutile (1 1 0) surfaces. The transition metal oxide surfaces considered were IrO2, RuO2, PtO2, and TiO2. It was found that the bridge sites were covered by oxygen for most conditions, so that the CUSs were active in OER while, for example, in the case of RuO2, O-covered Ru CUSs were active for ClER. The energy of adsorption of the different intermediates (Clc, OHc, Oc, Oc2, Cl(Oc)2, and ClOc) were then calculated, and a number of Pourbaix diagrams were determined. The Pourbaix diagram for RuO2 is reproduced in Figure 21. In each area of the Pourbaix diagram, the predominant species is the one which has the most negative free energy of adsorption versus a suitable reference state (e.g., gas phase Cl2 in the case of the Cl adsorbate). The Pourbaix diagrams indicated under which conditions chlorine or oxygen evolution would be more probable. Furthermore, as was the case for oxygen evolution,184 it was found that the binding energy of the intermediates (including the Cl-containing species) all scaled with ΔE(Oc) (the adsorption energy of Oc). Application of the scaling relations allowed a general phase diagram, where U and ΔE(Oc) decide which surface state is most stable, to be constructed (see Figure 22). This phase diagram is correct for all pure rutile oxide (110) surfaces, as long as all relevant intermediates have been considered. The diagram was constructed for pH = 0 and aCl− = 1, to make the diagram relevant at conditions where chlorine gas forms, but it can be adapted for other conditions as well. Only the areas which allow formation of stable chloride-containing intermediates are likely as chlorine evolution sites, since their stability likely makes the sites abundant on a real electrocatalyst. 3005


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Cl(Oc )2 + Cl−(aq) → 2Oc + Cl 2(g) + e− c

(41) c

for the Cl(O )2 intermediate. In these reactions, signifies a bare CUS and Occ 2 is an adsorbate with molecular oxygen-like geometry bound to two neighboring CUSs. For each reaction, an expression for the electrode potential where all steps are spontaneous, as a function of ΔE(Oc), can be written and entered into the general Pourbaix diagram (Figure 22). The resulting diagram is shown in Figure 23, which contains the

Figure 21. Phase diagram calculated by Hansen et al.183 The only area (at the pH and potentials included in the diagram) where the surface has a stable Cl-containing intermediate is at U > 1.55 V at pH < 1.3. Since the only stable chlorine intermediate is Cl(Oc)2, Hansen et al. believe that the reaction mechanism that yields Cl2 on real RuO2 surfaces includes the Cl(Oc)2 intermediate. Reproduced with permission from ref 183. Copyright 2010 PCCP Owner Societies.

Figure 23. General phase diagram, with thick lines indicating the “volcano” for chlorine evolution activity, taken from Hansen et al.183 Actual oxide materials are indicated as red dots with error bars. Dotted lines are the lines for the chlorine evolution reaction involving each chloride surface species (ClOc, Cl(Oc)2, or Clc). The volcano is limited by the reaction that requires the lowest applied potential to occur, for each ΔE(Oc). It is also limited by predominance areas where chloride containing compounds are stable, and therefore common, on the surface. The horizontal line is the equilibrium potential for chlorine evolution. It is seen that RuO2 has an activity that is close to the optimal one in the volcano plot. The blue dashed line is the line for oxygen evolution. It is seen that the potential necessary for chlorine evolution is always lower than the potential necessary for oxygen evolution. Reproduced with permission from ref 183. Copyright 2010 PCCP Owner Societies.

volcano plots. The volcanoes are limited both by the areas where chloride-containing intermediates are prevalent and where the chlorine evolution reaction involving these intermediates is spontaneous. Hansen et al. conclude their work by supposing, based on Figure 23, that there is room to find an oxide material which has an optimal ΔE(Oc), at which the ClER would progress with a lower applied potential and with a higher selectivity for chlorine evolution. No similar analysis at conditions for chlorate production has been published so far. In 2014, Exner et al.207 performed a complete reevaluation of the theoretical Pourbaix diagram for chlor-oxy species on the (110) surface of RuO2. All relevant intermediates were considered. The Pourbaix diagram produced, shown in Figure 24, mostly agreed with that of Hansen et al.,183 even though Exner et al. performed their study using the PBE exchangecorrelation functional instead of RPBE. However, a key difference was that Exner et al.207 did not find that the Occ 2 (oxygen molecule-like) adsorbate was the most stable adsorbate on the surface of RuO2 at applied potentials and pH below those required for ClER or OER, instead identifying 2 Oc as the most stable configuration. Still, the intermediates present at pH and potentials required for both ClER and OER on RuO2 were found to be the same in both the study of Exner et al.207 and the study of Hansen et al.183 Furthermore, the difference in stability between c the Occ 2 and 2 O was within the errors expected from GGA-level DFT calculations. Nevertheless, the results of Exner et al.207 agreed with the GGA-level DFT study of Wang et al.,208 who also

Figure 22. General phase diagram, with the dominant surface compounds indicated for each predominance area, of Hansen et al.183 The background is a visual representation of the surface and the compounds, where metal ions are blue, oxygen is red, hydrogen is white and chlorine is green. Reproduced with permission from ref 183. Copyright 2010 PCCP Owner Societies.

Three reaction mechanisms, one for each of the stable surface chloride compounds (ClOc, Cl(Oc)2, or Clc), were proposed based on the calculated adsorption energies. The effects of applied potential, pH, chloride concentration, and electrode material were considered, using the computational hydrogen electrode (CHE) model,195 when determining under which conditions each of the three reaction mechanisms would become spontaneous. The reaction mechanisms are the following Oc + Cl−(aq) → ClOc + e−

(36)

ClOc + Cl−(aq) → Oc + Cl 2(g) + e−

(37)

c

for the ClO intermediate, Cl−(aq) +c → Clc + e− −

c

(38) c

Cl + Cl (aq) → Cl 2(g) + +e

−

(39)

c

for the Cl intermediate, and c − − Occ 2 + Cl (aq) → Cl(O )2 + e

(40) 3006


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Exner et al.210 continued the theoretical study the year after, by considering the effect on the chlorine evolution activity of applying overlayers of other metal oxides on RuO2 (110). A significant finding was that deposition of a monolayer of PtO2 on RuO2 (110) should yield increased activity for ClER. Additionally, in 2015, Exner et al.211 reported results on the effects of doping another transition metal (either Cr, Ir, Mn, V, Au, Rh, Co, or Pd) in either the active CUS or the cationic bridge site next to the active site in RuO2. It was found that either doping Ir into the active site, or doping Cr, Ir, Mn, Zr, or V into the neighboring cationic site, could improve the activity for OER. However, in neither study was the selectivity between OER and ClER considered. Later in 2014, two theoretical studies of chlorine evolution selectivity, a study focusing on RTO by Karlsson et al.64 and a study focusing on monolayers of TiO2 on RuO2, by Exner et al.,63 were published. The study of Karlsson et al.64 was based on the previous volcano plot of Hansen et al.183 The adsorption energy of atomic oxygen on the CUS of TiO2, RuO2, and both Ru-doped TiO2 and Ti-doped RuO2, as well as on a structure with the same overall stoichiometry as that of conventional DSA (30% RuO2 and 70% TiO2) were studied (see Figure 25). The adsorption energy of oxygen had been found by Hansen et al.183 to serve as a descriptor for both OER and ClER on rutile oxides. The approach used by Karlsson et al.64 allowed the effect of the local structure of an electrode (i.e., the position of the dopant relative to the surface-active CUS) on the electrocatalytic properties of the mixed oxide to be studied. It was found that while doping RuO2 with Ti has a relatively minor effect, doping TiO2 with Ru serves to activate Ti CUSs. These activated Ti sites had oxygen adsorption descriptor values in the range optimal for highly selective and active chlorine evolution. If the dopant Ru atom was situated in the CUS, the descriptor value was the same as for pure RuO2, indicating that the properties of Ru were maintained even in the mixed oxide. These results are shown in Figure 26. The study indicated that the improvement in ClER selectivity that RTO electrodes exhibit, when compared with electrodes of pure RuO2, is related to activation of surface Ti atoms by nearby Ru dopant atoms. The study also indicated that even electrodes with monolayers of TiO2 on RuO2 should exhibit improved selectivity

Figure 24. Theoretical Pourbaix diagram for OER and ClER on RuO2 (110) according to Exner et al.207 (excluding solvation effects). The areas in which ClER and OER can occur are marked in the figure. The conditions (aCl− = 1 and T = 298 K) are the same as those in Figure 21, the Pourbaix diagram according to Hansen et al.183 The diagram is similar to that of Hansen et al.183 in most respects, but the key difference resides in Exner et al.207 finding that Occ 2 is unstable on the surface and that both CUSs are covered by Oc at potentials and pH values below those where ClER and OER can occur. Reprinted with permission from ref 207. Copyright 2014 Elsevier.

found that the Occ 2 adsorbate on RuO2 should dissociate into 2 Oc. The authors also took a step further by including the effect of solvation, by use of a continuum solvent model. This had a clear effect on the Pourbaix diagram and also made the 2 Oc configuration significantly more stable than the Occ 2 configuration. However, the effects of solvation on the adsorption energies of the different intermediates were found to be much larger than what was found in the first-principles studies of both Rossmeisl et al.184 and a later more thorough study by Siahrostami and Vojvodic.209 The latter calculation, in which water was considered explicitly, found that the effect of up to 2.5 monolayers water molecules on the adsorption energies of the OER intermediates (HOc, Oc, and HOOc) is minor for RuO2 (110).209 However, the effects of water molecules on the adsorption energies of chlorine species was not considered, warranting future study of the effects of solvation on such species.

Figure 25. Model structures considered by Karlsson et al.64 (a) is a side view of the structure of Ru-doped TiO2, (b) indicates is a side-view of the structure of the DSA model system, and (c) is a three-dimensional image of the surface formed by the system indicated in (a). In (a and b), the surface normal is in the z direction, upward in the figure. Reprinted with permission from ref 64. Copyright 2014 Elsevier. 3007


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Figure 26. Trends in oxygen adsorption energy descriptor values as the position of the dopant atom is changed, indicating an activation of Ti sites by dopant Ru atoms. This could explain the improved selectivity for ClER that RTO electrodes exhibit. Reprinted with permission from ref 64. Copyright 2014 Elsevier.

Figure 27. Results of doping rutile TiO2 with other dopants, in either the CUS (left side) or the bridge site (right side). The vertical lines correspond to descriptor values for pure TiO2 (thin green), optimal selectivity and activity for ClER (thick red), maximum OER activity (thick blue), and for pure RuO2 (black). Reprinted with permission from ref 65. Copyright 2015 Elsevier.

for ClER, while maintaining a similarly high activity as pure RuO2. Interestingly, the study of Exner et al.,63 building upon the results from their previous study,207 also concluded that a monolayer of TiO2 on RuO2 should optimize the selectivity for ClER. Thus, the two studies were in basic agreement in spite of being based on two independent sets of calculated surface

Pourbaix diagrams and volcano plots and used two different exchange-correlation functionals and assumptions related to effect of the electrolyte on the reaction. The reason for the improvement in selectivity is the decrease in Oc adsorption energy that occurs when applying a monolayer of TiO2 on RuO2.63,64 3008


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Figure 28. Pourbaix diagram for Ru in concentrated chloride solution, from Loučka.214 Original caption reads: “Potential-pH diagram for the Ru− H2O−Cl−, system at 25 °C in a solution of 264 g/dm3 NaCl. Activities of other ions are 1 × 10−6 M”. Reprinted with permission from ref 214. Copyright 1990 Springer Science and Business Media.

In a publication from 2015, Karlsson et al.65 expanded their theoretical studies of doped rutile TiO2.65 The method was analogous to the one used in Karlsson et al.,64 but now included 38 dopants, mainly transition metals. A general study of the effect of the dopant either as active site, or located close to a Ti site, was performed, suggesting a number of possible dopants that would activate TiO2 either for OER or ClER. The results are shown in Figure 27. Arsenic was suggested as a dopant that would result in maximal ClER activity regardless of position. The dopants Bi, Co, Ir, Mn, Pd, Ru, or V were predicted to result in a doped oxide with optimal ClER selectivity and activity, while doping with Mo or Re was predicted to result in an oxide with high OER activity. However, it was found that the local structure of the active site is key to achieving these effects. The properties of some ternary oxides were also studied, finding that the combined effect of two dopants located next to a Ti-active site was close to the average effect of the dopant in their respective binary Ti oxides. Furthermore, it was found that under OER conditions, but not under ClER conditions, there is a driving force for most dopants to segregate toward the active CUS. One of the highest driving forces was noted for Ru-doped TiO2, indicating that the deactivation of Ru-based DSAs might occur through segregation of Ru toward the active site, followed by formation of RuO4.212 However, the process after segregation toward the active site (e.g., formation of RuO4) was not studied in the paper.

works that have clarified the relationship between selectivity for Cl2 or ClO−3 and O2 and the stability of the electrode material will be considered. As will be seen, the decomposition of RTO electrodes is connected with the removal of Ru from the electrode in reactions related to the OER. For this reason, studies of both Ru metal, RuO2, and mixed Ru oxides will be considered in this section. The discussion will start from the Pourbaix diagram for Ru, which was published in 1966,213 and the review will then discuss papers published from 1966 and onward. The Pourbaix diagram for Ru213 indicates the regions of stability of the pure metal (see Figure 28 for the Pourbaix diagram for Ru in chloride solutions of Loučka214). The conversion from Ru, to Ru(OH)3, RuO2, and finally volatile RuO4 at potentials above 1.5 V versus SHE at pH 0 is central to the understanding of the stability of Ru and its oxides. The potential required for conversion of RuO2 to form RuO4 decreases with increasing pH, until the final product of the oxidation of RuO2 becomes RuO−4 and then RuO2− 4 at pH > 9− 12. Around the same time, Llopis and Vázquez215 and Llopis et 216 al. studied the corrosion of Ru. In the first study, anodic and cathodic charging curves for Ru in HClO4 and HCl solutions were measured. On the basis of these results, the electrode potentials where oxidation of Ru to form Ru2O3, RuO2, and RuO4 occurs were found. In Llopis et al.216 a radiochemical technique, which allowed for the detection of very small amounts of dissolved Ru, was applied to study the potentiostatic corrosion of Ru in 2 or 4 M HCl, at 50 °C. The technique was based on detecting a radioactive Ru isotope, which is formed after neutron irradiation of the electrodes in a nuclear reactor. The potential in the electrochemical measurements was kept constant at a

2.3. Electrochemical Studies Concerning the Relationship between Selectivity and Stability

There are many studies of Ru metal oxide degradation and decomposition. A complete discussion of this literature would necessitate a full review. Therefore, in this section a selection of 3009


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that Ru oxides form on the surface of the electrode during OER and that ruthenates are released into the solution. In 1978, Uzbekov et al.221 presented a radiochemical study of decomposition of RTO electrodes prepared using thermal decomposition to achieve the standard composition of 30% RuO2 and 70% TiO2. The electrodes were polarized anodically in 5 M NaCl, kept at pH 7, at several potentials during 10 to 80 h. Initially, measurements were conducted at an anodic current density of 2 kA/m2 and an anode potential of 1.34 V versus SHE. The authors found that the corrosion rate, as measured by the radioactivity of dissolved Ru in the solution measured during the experiment, decreased from more than 10 × 10−6 g/cm2 h to around 10 × 10−7 g/cm2 h over the first few hours and then steadily decreased further over time. Although the solution was continually stirred, the radioactivity measured when draining the cell after the end of the experiment was higher than that measured in the circulating solution. This indicated to the authors that a part (10 to 40%) of the degradation of the anode took place by detachment of solid particles from the surface. Ru was also dissolved in the electrolyte, indicated by the deposition of Ru on the counter electrodes and released into the gas phase as RuO4. It was estimated that 5−15% of the Ru was released into solution and that 10% was released into the gas phase. Furthermore, it was noted that the corrosion rate increased 2− 5 times upon current interruption of even a few seconds. After several hours of electrolysis, the corrosion rate again decreased to a low value. Finally, the authors examined corrosion rates at several different anodic potentials. A significant increase in the corrosion rate was noted at potentials above 1.3 V versus SHE. Increasing the anode potential increased the rate of corrosion, but the specific rate (per mole Cl2 produced) decreased. In the same year, Kokoulina et al.222 presented polarization measurements on conventional RTO anodes measured in 1 N H2SO4 at different pH in the range of 0.5−13 and temperatures between 20 and 80 °C. An increase in the Tafel slope was noted at 1.5 V versus RHE at 20 °C and 1.4 V versus RHE at 80 °C. The potential at which this bend occurs was termed the critical potential (a term first introduced by Veselovskaya et al.219) at which the change in Tafel slope might signal an increase in corrosion rate due to formation of RuO4. Increasing the pH decreased the critical potential, which would agree with the Pourbaix diagram.213 Also in 1978, results based on anodic polarization measurements conducted in buffered sulfuric acid solutions using pure RuO2/Ti electrodes was presented by Bondar and Kalinovskii.223 The authors determined the critical current density (the current density at the critical potential,219,222 at which the change in Tafel slope of the polarization curve occurs) as a function of pH at 20 and 60 °C. This indicated the current densities which should not be exceeded for RuO2 electrodes, at various pH values. As an example, the critical current density for oxygen evolution on RuO2 in sulfuric acid at pH 0 was approximately 1 kA/m2. In 1979, Gorodetskii et al.224 examined the effect of current density on the corrosion rate of DSA. Again, electrodes with the standard 30:70 Ru:Ti composition were prepared by thermal decomposition. Electrolysis was performed at 87 °C in 5 M NaCl at a pH of around 2, and the corrosion rate of Ru was determined using a radiochemical method. Results for one measurement carried out at a current density of 2 kA/m2, giving an anode potential of 1.33 V versus SHE, was shown. As in Uzbekov et al.,221 the corrosion rate decreased slowly over several hours and increased for a short time after shutdowns when the electrode had been removed from the solution. It was estimated that 14%

number of potentials in the range from 0.95 to 1.25 V versus SCE during more than 10 h. The Coulombic corrosion yield (the percentage of the current resulting in the corrosion of Ru) decreased sharply as the potential was increased above 1.0 V versus SCE. At these potentials, a passivating Ru surface oxide grows on the metal, decreasing the rate of Ru cation transport through the film. At potentials approaching 1.4 V, formation of RuO4 started.213,216 In 1973 and 1976, Gorodetskii et al.217 and Pecherskii et al.,218 respectively, presented further work on the anodic behavior of Ru. The studies examined the anodic dissolution behavior of Ru in the potential range from 0.8 to 1.4 V versus SHE. A radiochemical technique was again used to determine the amount of dissolved Ru. In the first study, the anodic behavior of the electrode was studied in both 3 M HClO4, where oxygen evolution is the dominating anodic process, and in 3 M HCl, where chlorine evolution is the dominating anodic process. In the former solution, the total current density as well as the rate of Ru dissolution increased as the potential was increased above 1.2 V versus SHE. The slopes for both processes were equal, at about 30 mV per decade in either corrosion rate, measured in g/cm2 s or current density. The authors concluded that the rate of the OER was connected with the rate of Ru dissolution, possibly due to the existence of shared surface intermediates. However, in 3 M HCl at the same potentials, although the rate of Ru dissolution did increase with increasing electrode potential, the slope of the Ru dissolution rate was significantly lower than that of the total current density. It appeared that during chlorine evolution, increasing current density resulted in a decreasing current efficiency for Ru dissolution. Furthermore, Gorodetskii et al. studied the effect of increasing Cl− concentration from 0.01 to 1 M at an anode potential of 1.4 V versus SHE, finding a 4-fold decrease in the Ru dissolution rate together with an increase in overall current density by a factor of 300. The authors suggested that electrodes, specifically RTO, with increased selectivity for the ClER thus should also exhibit improved corrosion resistance. Pecherskii et al.218 performed a deeper study of the effect of pH on the Ru dissolution rate. The overall anodic current density and the Ru dissolution rate were determined in solutions of HClO4, H2SO4, or HCl at a constant ionic strength of 3 M and acid concentrations from 0.01 to 3 M. In the sulfate and perchlorate solutions, both the overall current density and the Ru dissolution rate increased with increasing pH, again indicating the connection between Ru dissolution and the OER. In the HCl solution, the overall current density decreased slightly with increasing pH, and the Ru dissolution rate increased. In 1974, Veselovskaya et al.219 published a study of the anodic behavior of metallic Ru including measurements conducted in a number of solutions, including 300 g/dm3 NaCl at pH 3. All measurements were performed at 80 °C. The authors noted that the regions of Ru dissolution at lower anodic potentials (below 1 V vs SHE) are followed by regions of lower Ru dissolution rate (due to the formation of hydrous RuO2) at potentials above 1 V versus SHE. The Tafel slope for electrolysis in the chloride solution at these higher potentials coincided with the typical 40 mV/dec Tafel slope for ClER on a DSA electrode, supporting the idea that a hydrous RuO2 layer covers the surface under these conditions. The paper is perhaps most notable due to its introduction of the “critical anode potential”, Ecr, here given as approximately 1.45 V versus SHE, as the potential where conversion of RuO2 into volatile RuO4 becomes favorable. In 1977, Iwakura et al.220 presented research on possible mechanisms for OER and for Ru dissolution from Ru, suggesting 3010


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of particles, independent of the pH value of the solution. The results indicated that the connection between the OER and the rate of corrosion is not necessarily as simple as was previously indicated. At pH below zero, the corrosion rate was said to decrease significantly when increasing the pH up to 0. In the pH range from 0 to 3, the rate of corrosion of the DSA was essentially constant at a value of about 1 × 10−8 g/cm2 h in the concentrated NaCl solution. The corrosion rate was somewhat higher in the 100 g/dm3 NaCl solution, here showing a minimum at about pH 1, and then increased with increasing pH for both concentrations. However, the OER rate showed a clear increase when increasing the pH from 0 to 2, and then again when increasing the pH from 4 to 5, in the 300 g/dm3 NaCl solution. Increases in the OER rate was also noted when increasing the pH from 2 to 4 in the 100 g/ dm3 NaCl solution. The results thus indicate a region of maximal stability around pH 0 to 3 in the 300 g/dm3 NaCl solution at industrially relevant current densities. Furthermore, the total percentage of Ru released as gaseous RuO4 was determined as a function of pH at both NaCl concentrations. In both cases, at a

of the corrosion was due to detachment of particles and that 12% was due to formation of gas phase RuO4 (apparently being produced at an anode potential below 1.4 V vs SHE). A comparison between the corrosion rate of pure RuO2 and of the RTO electrodes indicated that the corrosion rate of the pure Ru oxide was 1 order of magnitude higher than that of the RTO material. The total corrosion rate of the RTO anode was also shown as a function of the current density, once more indicating that the rate of RTO degradation increases with increasing current density but that the specific (per mol Cl2 produced) corrosion rate decreased with increasing current density. In 1980, Barral et al.225 published results on the deactivation of RuO2 and RTO during OER in 1 M H2SO4 at 40 °C. The deactivation was determined indirectly, by measuring the time required to reach an anodic potential of 4 V versus the saturated mercury sulfate electrode. Interestingly, it was found that pure RuO2 electrodes had a longer lifetime than RTO electrodes at 10 kA/m2. However, the electrodes had been prepared by thermal oxidation at 500 °C for only 10 min with no final prolonged heating. This is significantly shorter than the conventional time of about 1 h for the final oxidation of the coating. Indeed, Gerrard and Steele11 note that for pure RuO2, about 5 min is needed to achieve a stable value for the electrical conductivity, but an oxide coating of 30:70 Ru:Ti needed more than 1 h to reach a stable conductivity. Therefore, it is possible that the results of Barral et al.225 disagree with the majority of other work since their mixed oxide coatings were not sufficiently stabilized during preparation. Still, they noted a trend for the pH dependence of the deactivation of pure RuO2 that seems to contradict that of Bondar and Kalinovskii.223 Barral et al. found that the lifetime was essentially constant in the range of pH from 0 to 13 at a current density of 5 kA/m2. At higher pH, the lifetime increased significantly. The authors suggested that the reason for the constancy in lifetime is related to the difference between bulk and anode surface pH. Because of the ongoing OER, the pH at the anode is lower than in the bulk. Only at very large pH values (e.g.,

−

current density of 2 kAm 2 and anode potential of 1.33 V versus SHE, the percentage released as RuO4 increased significantly when increasing the pH, from about 10% at pH 0−1 to about 60% at pH 4 in the 100 g/dm3 NaCl solution and 40% at pH 5 in the 300 g/dm3 NaCl solution. The work of Pecherskii et al.67 from 1982 has already been described in section 2.2.3, but the corrosion results will here be discussed in more detail. Again, the study in question examined the properties of three types of electrodeposited coatings on Ti substrates. The first type consisted of Ru only, the second of Ru which was then heat-treated for several hours, and a third which was treated the same way as the second but was then polarized anodically in sulfate solution until the electrode potential increased sharply. The first electrode could thus be considered to consist of a Ru hydroxide or metal, the second of RuO2, and the third of RuO2 mixed with TiO2. Auger electron spectroscopy was performed, indicating that Ti was present in the surface layer of the third type of electrode but not in the second type. The same trend as was noted for the oxygen production rate was seen also for the corrosion rate. Electrodes of the third type not only produced less oxygen but also exhibited significantly lower corrosion rates than the second type. The corrosion rate at 1 kA/ m2 was about 1 × 10−5 g/cm2 h for the first type of electrode, about 1 × 10−6 g/cm2 h for the second type of electrode, and about 1 × 10−9 g/cm2 h for the third type of electrode. The pH dependence of the corrosion was also studied. Electrodes of the third type exhibited the behavior seen by Gorodetskii et al.,151 of essentially constant corrosion rate up to pH 3, followed by an increase in corrosion rate. Corrosion rates increased with anode potential at potentials positive of 1.3 V versus SHE for all three types of electrode. The slopes of the increase in corrosion rate with electrode potential was similar to the Tafel slope of the oxygen evolution reaction. In 1982, Burke et al.228 presented results for OER using RuO2. As in the study of Barral et al.,225 the lifetime was estimated based only on a steep increase in anode potential. Nevertheless, it was again seen that increasing current density results in a significant decrease in lifetime in both acid and base. The lifetime of the anode decreased significantly with acid concentration in H2SO4 solutions, while in alkaline solution of NaOH the lifetime increased significantly with increase in concentration above 1 M NaOH. Mechanisms for both OER and Ru corrosion were suggested.

−

a bulk pH > 11.5 at a current density of 100 Am 2), does this change, they argue. Nevertheless, the fact that the corrosion rate was not determined directly by measurement of the corrosion products limits the utility of the study. Determining the lifetime based only on an increase in anode potential means that the deciding factor for the lifetime might be the formation of a nonconductive interlayer between the Ti substrate and the coating, rather than decomposition of the coating itself. It is therefore difficult to disentangle the various mechanisms of deactivation when using only the potential as the deciding factor. Still, it will be indicated much later in the works of Tilak et al.226 and Eberil’ et al.227 that the appearance of an insulating TiO2 layer might not be an important mechanism for DSA deactivation. In 1981, Gorodetskii et al.151 published a study of the effect of pH in the range from 0 to 5 on both the corrosion rate and on the rate of the OER on DSA. Some corrosion measurements were also carried out at lower pH values. Again, standard DSA electrodes were studied at 87 °C in 300 g/dm3 or 100 g/dm3 NaCl. Once more, electrolysis was carried out at a current density of 2 kA/m2. The same radiochemical method as used previously was employed. However, the O2 formed was now quantified by GC. Corrosion measurements were started after the system had reached steady state (after about 200 h), and measurements with different pH and NaCl concentration were then performed. Again, about 14% of the corrosion was the result of detachment 3011


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current density of 2 kA/m2 at 60 °C. Corrosion rate measurements were conducted using the same radiochemical method as previously applied. The effect on the anode potential and corrosion rate of changing the electrolyte concentration from 5 M NaCl to 6 M NaClO3 was examined, finding a clear increase in anode potential as the NaCl concentration was decreased. At the same time, the corrosion rate increased. The mechanism involved is likely the same as has been seen in several other studies, that the rate of OER increases significantly when the Cl− concentration decreases, resulting in an increased corrosion rate. In 1986, Kötz and Stucki232 presented results on stability of mixed Ir−Ru rutile oxide coatings. While the corrosion rate was only determined based on the standard accelerated lifetime test (as determined by an increase in the anode potential) in H2SO4, the authors indicated that mixing RuO2 and IrO2 results in stabilization of the material, although at the cost of a decrease in activity for OER. In the same year, Kolotyrkin233 presented a review of the relationship between corrosion and kinetics for anodes. For RuO2 and DSA, the results essentially mirror the discussion thus far: for RuO2 both the rate of OER and rate of corrosion decrease in a similar way as a chloride solution is acidified. For DSA, the dissolution rate exhibits a minimum in the pH range 1−3, as seen by Gorodetskii et al.151 and Uzbekov and Klement’eva.229 In 1987, Lyons and Burke234 presented a study of the stabilization of RuO2 by addition of SnO2 in which the lifetime was evaluated by a standard accelerated lifetime test during OER in concentrated NaOH at 80 °C. In 1989, Zhinkin et al.163 presented results on the corrosion behavior of standard DSAs, focusing on current densities exceeding 2 kA/m2 and several temperatures in the range from 70 to 90 °C. The pH range was 2 to 5. Measurements were conducted in 300 g/dm3 NaCl. The same radiochemical method as had been used in several previous Soviet studies was again applied. The authors analyzed the gas phase and solid phase particulate losses of Ru, finding that the gas phase loss made up 20−40% of the total loss and that the solid particulate losses were less than 1% of the total corrosion. A clear increase in the corrosion rate was noted as the pH was increased above 4. This increase was concomitant with a clear increase in the percentage of O2 in the off-gas as the pH was increased. The connection between Ru corrosion and the rate of the OER was thus again seen. Increasing the current density from 5 to 10 kA/m2 also increased the corrosion rate but less significantly. There was a clear effect of temperature on the corrosion rate, with higher corrosion rates noted at 90 °C than at 70 °C. Still, the effect of increasing the pH from 2 to 5 was more significant than both the effect of current density and of temperature. In 1990, Manli and Yanxi235 presented results on corrosion of DSA, prepared using the conventional method, during electrolysis at 10 kA/m2 in saturated NaCl solution at 80 °C. The corrosion rate was evaluated based on when the electrode passivated, resulting in a high anode potential. Secondary ion mass spectrometry (SIMS) results indicated that Ru is being selectively removed from the surface layer of the electrode during electrolysis. This agrees with the results found by Uzbekov and Klement’eva.229 In 1990, Loučka214 presented a Pourbaix diagram for Ru in concentrated chloride solution, as relevant for chlor-alkali production. The main detail important to notice from this Pourbaix diagram is the supposed existence of soluble Ru chlorides at low pH and anodic potentials. Interestingly, the

In 1983, Denton et al.154 presented results from radiochemical experiments using RTO electrodes used for chlorine evolution in either membrane or diaphragm cells, finding a direct relationship between the rate of Ru loss from the coating and the oxygen percentage in the cell gas. It was also claimed that increased pH of the anolyte resulted in increased anode wear, that increased cell temperature or current density resulted in increased anode wear, and that increased chloride concentration was associated with a decrease in anode wear rate. In 1984, Kötz et al.212 contributed a detailed study of the corrosion of RuO2, in which CV was combined with in situ differential reflectance spectroscopy, finding that during OER in H2SO4, RuO4 is the only corrosion product. Furthermore, the release of RuO4 started at the same potential as the OER. This study clearly shows the mechanistic link between the OER and the corrosion of RuO2. In 1985, Uzbekov and Klement’eva229 published a study employing radioactive isotopes of Ru and Ti in the salts used to prepared a DSA coating, so that both Ru and Ti corrosion products (regardless of oxidation state) could be detected during electrolysis. The experiments were conventional chlorine evolution experiments, conducted in 5 M NaCl and 90 °C, with a current density of 2 kA/m2. The effect of pH, in the range from 0 to 11, was studied. The electrode potential was essentially constant at about 1.33 to 1.38 V versus SHE in all experiments, indicating the constancy of the activity of the electrode during corrosion, over time, and at all pH values. It was found that there are clear differences in the dissolution of Ru and Ti from RTO electrodes. During a long-term electrolysis experiment at pH 2, during the first 30 h, removal of Ru from the coating was the dominating process. Although the ratio between Ru and Ti were 0.45 to 1 in the coating, during the first 30 h, the selectivity for dissolution of Ru over Ti decreased from values between 4 and 2 toward equal dissolution. After 30 h, the selectivity was close to 1, indicating equal rates of Ru and Ti dissolution. This means that during electrolysis, the surface layers of RTO lose Ru and become more concentrated in Ti. Measurements of the steady state corrosion rate of Ti and Ru as a function of pH were also performed. At low pH (around and just under 0), the corrosion rates of Ru and Ti were both elevated, but the selectivity for Ru dissolution was less than 1, indicating a higher rate of Ti dissolution. As the pH then was increased, the dissolution rate of Ti decreased, reaching stable values above 2 and up to a pH of 5.4. One measurement was apparently then made at pH 11.3, at which the rate of dissolution of Ti had increased significantly. The dissolution rate of Ru showed a minimum at about pH 2 then increased. The resulting trend in the selectivity for dissolution of RuIV versus TiIV is an increase from less than 1 at the lowest pH, increasing to almost 9 at pH 5.4. Due to the clear increase in Ti removal rate at pH 11, the selectivity decreased to 4. Aside from the lowest pH values below pH 2, the dissolution rate of Ru was up to 9 times higher than the dissolution rate of Ti. It was proposed that both Ti and Ru dissolve due to formation of soluble metal chlorides at the lowest pH values, while Ru removal could also occur due to the electrochemical conversion of RuO2 to RuO4 at higher pH values. At high pH values, Ti would be converted into soluble Ti oxyhydroxides, in agreement with the Pourbaix diagram.230 Further work from the same group appeared in 1985, this time focusing on the corrosion of RTO during sodium chlorate production.231 DSA were used, prepared using the conventional thermochemical method and with the standard 30:70 Ru:Ti composition. Measurements were conducted using an anodic 3012


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In 2003, Cornell et al.174 studied the critical anode potential using commercial DSA for chlorate electrolysis. It is known from industrial practice that during the first months of industrial operation in a chlorate cell, the anode potential decreases simultaneously with an increase in chlorate selectivity. From laboratory studies, it was concluded that a new, relatively compact DSA operates at a potential higher than Ecr. When in production, the real surface area of the electrode increases (probably due to corrosion of the coating) and the electrode potential thus decreases to a stable value at, or below, Ecr. The critical current density, at which the anode operates at Ecr, depends on factors such as the chloride concentration. Later work by Nylén and Cornell238 concluded that process conditions in industrial chlorate production leads to a higher risk of exceeding Ecr than the conditions in the chlor-alkali process, where the chloride concentration is higher, pH lower, operating temperature higher, and there is no chromate in the electrolyte. In 2004, Gajić-Krstajić et al.239 presented the application of a spectrophotometric method to determine low concentrations of dissolved Ru species on the corrosion of RuO2 deposited on Ti. The method was based on the catalytic action of Ru on the As− Ce redox reaction in sulfuric acid. This method is of a comparable sensitivity (1 × 10−9 M Ru) as the radiochemical methods employed in several studies mentioned previously in this section. The experiments were conducted in 0.5 M H2SO4 solution, at a current density of 5 kA/m2. The service life, determined in the standard way, increased exponentially with decreasing applied current density (and thus with decreasing rate of OER). By performing measurements at several potentials in the range from 1.52 to 1.67 V versus SCE, a constant dependence of the current density for oxygen evolution, jO2, on the current density for Ru dissolution, jRu, was found

diagram of Loučka only includes a very small pH range of RuO2 stability above the Cl2/Cl− equilibrium potential, between pH 2.3 and 3.0. It was pointed out that this pH range roughly agrees with that found to give the highest stability of DSA electrodes.151 In accordance with the Pourbaix diagram of Loučka,214 at lower pH, formation of soluble Ru chlorides was predicted to become favorable, while at higher pH, formation of (per)ruthenates becomes favorable. In 1998, Hardee and Kus236 reported on studies of passivation of DSAs through formation of an interlayer between the coating and the substrate. Electrodes deactivated either in the industrial setting or by using accelerated aging regimes exhibited changes in impedance phase angle at low frequencies (around 1 Hz) as well as in high-frequency regions (around 1000 Hz). These changes were attributed to formation of a TiO2 interlayer, although no direct detection of this interlayer was presented. A DSA coating with a number of TiO2 layers deposited on the surface showed changes only in the low-frequency region. Using XPS and atomic emission spectroscopy (AES), removal of Ru in the outer layers of DSA after deactivation was noted. This was not deemed to be the reason for the deactivation, as such removal was also noted for electrodes that still functioned with lower anode potentials. Furthermore, a pure IrO2 electrode was deactivated in a sulfate electrolyte, after which changes particularly at high frequencies were noted. Also on the basis of this observation, it was concluded that formation of an interlayer was the mechanism of deactivation, as a significant amount of IrO2 still remained on the electrode. In 2001, Tilak et al.226 presented an electrochemical impedance spectroscopy (EIS study of the deactivation of DSA electrodes. An accelerated aging process was employed, wherein the potential was changed from 1.35 to −0.32 V versus SCE at 60 Hz. These results support those reported several times before, that during electrolysis, corrosion results in the selective removal of Ru from the surface layers of the electrode. Tafel, CV and EIS results indicated that deactivated electrodes (as determined by the standard passivation test of an increase in anode potential) behave similarly to fresh electrodes prepared with a 5% Ru content. Furthermore, the authors note that fresh electrodes with 5% Ru behave the same way in the low frequency region as aged 30% Ru electrodes, indicating that the deactivation process of DSA electrodes is due to removal of Ru from the surface, rather than formation of an insulating TiO2 layer between the Ti substrate and the coating. In the same year, Eberil’ et al.227 presented results on aging of DSA electrodes under standard chlorate production conditions. It was seen that the anode lifetime increased with the amount of Ru in the electrode, indicating the relationship between anode lifetime and the dissolution of Ru from the coating. Furthermore, as Tilak et al.226 did in the same year, Eberil’ et al. also performed tests to examine the proposed deactivation mechanism of formation of a TiO2 layer between the coating and support. It was seen that a slight improvement in anode lifetime resulted from the insertion of a Pt sublayer between the support and the coating. However, introducing Pt into the coating itself resulted in a significantly improved lifetime. This also indicates that the reason for DSA deactivation is the removal of Ru from the coating, rather than the formation of a nonconductive interlayer. Also in 2001, Bommaraju et al.237 presented results on deactivation of DSA electrodes. Much of the work was presented in Tilak et al.,226 but the authors also showed SIMS results agreeing with those of Manli and Yanxi.235

jRu jO

2

= 8 × 10−6 (42)

In 2014 and 2015, a number of papers of Mayrhofer and coworkers240−243 were published. In all papers, an in situ method to measure transition metal degradation products was applied, in which an electrochemical flow cell is coupled to an inductively coupled plasma-mass spectrometer (ICP-MS).244 First, Zeradjanin et al.242 reported results on the connection between RuO2−SnO2 mixed oxide surface morphology and OER activity and stability. Electrodes with two types of surface morphology, either cracked or smooth, were prepared using thermal oxidation of sol−gel-synthesized precursors onto Ti substrates. The usage of more dilute solutions of the Ru and Sn precursors in isopropanol, together with repeated coating application cycles, enable the formation of smooth surfaces. Electrolysis was performed in a 0.1 M H2SO4 electrolyte at 25 °C. The sensitivity of the ICP-MS instrument, allowing detection of nanograms of Ru per square centimeter of electrode area, is such that dissolution of Ru from the RuO2 component of the mixed coating during cycling of the potential in the range from 1 to 1.6 V versus RHE could be measured. For an industrial catalyst, the current density is normally kept constant, apart from at shutdown, so these results are not directly comparable to a steady-state dissolution rate during industrial electrolysis. Furthermore, it is unclear what the ratio between RuIV and SnIV components in the mixed coatings were. Still, it was found that the dissolution rate of Ru increased as the potential was increased over the reversible potential for OER of about 1.25 V versus RHE, and became especially high above 1.45 V versus 3013


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Table 4. Effects of Process Conditions on the Selectivity between Oxygen and Chlorine or Chlorate Formation Identified in Literaturea selectivity for oxygen evolution

factor

refs

comment

low pH (chlorine production) pH for chlorate conditions higher cCl− increased j higher cHOCl higher cNaClO3

decreased optimum pH value exists decreased decreased increased decreased

41,151,154, and 163 56, 67, 141, 162, and 173 66, 70, 75, 76, 121, 128, 154, 155, 162, 170, and 173 54, 56, 57, 62, 66, 69, 70, 75, 131, 154, and 162−164 54, 55, 57, 71, 116, 121, 130, 162, and 170 6, 54, 162, and 173

increased cell temperature anode potential (chlorate production) higher anode (changed geometry) larger holding volume irradiance

increased optimal anode potential exists decreased decreased increase

54, 57, 74, 154, 162, and 164 74 and 164 57 55 and 62 6 and 94

effect of pH is complex especially if cCl− < 100 g/dm3 increase at low cCl−170 related to viscosity?

current efficiency not directly measured

a The table focuses on trends described in concentrated NaCl solutions. This table gives a simplified view of the factors involved. The trends are described in more detail in section 3.1.1.

Table 5. Effects of the Electrolyte Composition on the Selectivity between Oxygen and Chlorine or Chlorate Formation Identified in Literaturea factor

selectivity for oxygen evolution

refs

comment/contradictory results

F− PO3− 4 SO2− 4 NO−3 Na2Cr2O7 Al Ag As Ce Co

possible decrease increase increase increase increase? no effect increase increase no effect increase

Cr Cu Fe

no effect increase increase?

137, 144, and 173 137, 157, and 173 138 and 157 137 137 peroxide,96 oxide92 oxide sources cited in ref 93 metal94 salt92 salt/ionic,86,92,94,97,100 metal,94 oxide93,95−97,105,106 and sources in,93 unspecified168 salt,92 oxide95 salt/ionic,90,97,103 metal,94 oxide,94,96,97,105 unspecified168 metal,94 oxide sources in93,96

Hg Ir Mn Mo Ni Pt Sn RTO, Pt/Ir Ru Rh U, W, Pd, Os, Tl, V Sb, Pb, Zn, Cd

increase? increase increase? no effect increase no effect increase? no effect no effect no effect no effect

oxide96 salt/ionic,92,101,102 oxide,95 unspecified168 oxide sources in93,96 oxide92 salt/ionic,103 metal,94 oxide95,96,105 and sources in93 unspecified168 unspecified168 metal94 particulates159 salt/ionic,92,95 oxide,92 unspecified168 salt/ionic95 oxide95

no effect

metal94

137

found no effect not with Pd71 137 found no effect no effect,92,95,162 slight decrease173 salt92 has no effect

less active than Co, Ni salt,86,92 metal,93 oxide92,95 or unspecified form168 has no effect

salt,86 metal93 has no effect

Ir increase according to ref 62

a

Several of the sources are studies of decomposition of bleach (sources cited in refs 93−95) or hypochlorous acid solutions (refs 90, 92, 93, 96, 97, 100, 102, 103, 105, and 106, which differ significantly from the electrolytes used in industrial chlor-alkali or sodium chlorate production. This table gives a simplified view of the factors involved. Sources cited in ref 93 are published before year 1900. Each factor is discussed in more detail in section 3.1.2.

thereafter, Cherevko et al.241 reported on the dissolution of IrO2 films, prepared using spin-coating onto Ti substrates and conventional thermal oxidation, during OER in 1 M HClO4. It was found that the stability of the films increased with increasing calcination temperature, in the range from 250 to 550 °C, except for electrodes prepared at 350 °C which was found to be the least stable. During constant-current experiments, performed at 20 A/ m2, it was found that dissolution of both Ir and Ti occurred, with

RHE, where the conversion of RuO2 to RuO4 can start. Dissolution was also noted at potentials below 1.23 V versus RHE, during cycling between 1.23 to 0 V versus RHE. It was found that the crack-free coating experienced a higher rate of decomposition in all potential ranges studied. Cherevko et al.240 then reported results on dissolution of metallic electrodes of Ru, Ir, Rh, Pd, Pt, and Au, finding a clear correlation between a low Tafel slope for OER and a high dissolution rate. Shortly 3014


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Table 6. Effects of the Anode Composition on the Selectivity between Oxygen and Chlorine or Chlorate Formation Identified in Literaturea factor

selectivity for oxygen evolution

modified composition

decrease

increased surface area larger coating crystallite particle size higher oxide loading less Ru in RTO coating

increase decrease

decreased j0 for reaction 11 IrO2 substitution Co substitution SnO2 substitution HfO2 substitution Zn substitution Ni substitution Mg substitution Fe substitution overlayer material support material

decrease decrease? decrease? decrease increase increases decrease decrease decrease? decrease possible effect

increase decrease

refs

comment

36, 158, 162, 167, 169, and 173 69 78 75 67, 68, 75, 139, 154, 161, and 171 69, 150, 159, and 167 74 78 69, 71, 147, and 172 69 181, 182, and 189 77 193 180 60 8 and 179

several modified coatings exist related to decrease in jloc? related to increase in jloc?

same effect as decreased j0 for reaction 11? Connected to increased chlorate formation36,141,150,154 but reaction 11 not only source doped into RTO; dorrosion products might increase selectivity in practice doped into RuO2;82,83,189 saw increase doped into RuO269,172 or together with Pd71 doped into RuO2 doped into RuO2 doped into RuO2 doped into RuO2 doped into RuO2, high Cl2 selectivity in Ar-saturated solutions overlayer of Pt mixed with IrO2

The table is limited to oxides based on RuO2. “Modified composition” signifies studies where details about the composition or physical properties of the electrodes were not given. Each effect is described in more detail in section 3.1.3. The table only includes studies where selectivity was measured directly [i.e., not only by comparing differences in electrochemical behavior (e.g. Tafel slopes) in solutions with and without chloride ions]. a

3. DISCUSSION

the dissolution rates of both components decreasing with time. In 2015, Reier et al.243 reported a combined OER activity and stability study of Ir−Ni oxides (rutile-like at Ni contents below 21% and of a layered brucite-like structure at higher Ni contents) in 0.1 M HClO4. It was found that coatings containing 67% Ni (as measured by inductively coupled plasma atomic emission spectroscopy) were most active for OER, which was related to a higher active surface area of this composition. However, the Ir dissolution rate of the coating with this composition was 34 times higher than the dissolution rate of pure IrO2, even at a low current density of 10 A/m2. Also in 2014, Frydendal et al.245 presented a study where electrochemical quartz crystal microbalance (EQCM) and ICPMS measurements were combined to monitor corrosion of rutile RuO2 and amorphous MnOx. The electrodes had been prepared using sputter deposition on Au or EQCM Au quartz crystals. Constant potential, at 1.8 V versus RHE for RuO2 and 1.8 and 1.9 V versus RHE for MnOx, and constant current, at 300 A/m2 for RuO2 and 200 A/m2 for MnOx, measurements were carried out in either 0.05 M H2SO4 in the case of RuO2 or 1 M KOH in the case of MnOx. Constant mass loss rates were noted in both types of measurement, which were carried out for 2 h. These electrodes thus did not exhibit any stabilization during electrolysis, in contrast with what has been found for both pure and mixed rutile oxides in other studies. A year later, Frydendal et al.246 found a that the stability of MnO2 coatings for OER in 0.05 M H2SO4 could be improved by codeposition of other transition metals. A combination of theoretical and experimental results indicated that mixed MnO2−TiO2 coatings, prepared by cosputtering, exhibited an improved stability at the cost of a slight decrease in activity.

3.1. Factors Affecting the Selectivity between Chlorine and Oxygen Evolution in the Production of Chlorine and Chlorate

The research literature portrays some general trends in the selectivity between chlorine and oxygen evolution. Some factors have been examined in several articles and can be considered well-understood. Others have been reported only once, or, in some cases, with different authors finding different results. Sections 3.1.1, 3.1.2, and 3.1.3 will discuss the selectivity from the point of view of process, electrolyte, and anode factors, while section 3.2 will discuss the relative importance of the different reactions that can lead to formation of oxygen. Tables 4, 5, and 6 summarize the different factors, related to process conditions, electrolyte contents, and electrode composition, respectively, that have been found in literature, as well as the literature references for each factor. The purpose of the discussion is to highlight the trends that have been reported, as well as to point to factors that have been reported to have opposing effects in different studies. The discussion focuses specifically on the trends that are relevant for chlor-alkali and sodium chlorate formation and is thus limited to concentrated NaCl solutions. In some cases, the trends in dilute solutions are different, as shown by, for example, the study of Czarnetzki and Janssen.170 Finally, the connection between anode stability and anode selectivity will be discussed in section 3.3. 3.1.1. Influence of Process Conditions. There are two factors that, at first, seem relatively straightforward to consider: the thermodynamics of the two main competing reactions 2 and 11 (written again below) and the effect of the current density on the selectivity. The equilibrium between chloride ions and chlorine is described by Cl− ⇌ 3015

1 Cl 2 + e− 2

(2) DOI: 10.1021/acs.chemrev.5b00389 Chem. Rev. 2016, 116, 2982−3028

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For the OER, the equilibrium can be written as H 2O ⇌

1 O2 + 2H+ + 2e− 2

ηp = A p + bp × log(j)

for the parasitic reaction. If the Tafel slope b for the main reaction is lower than the Tafel slope of the parasitic one, it is clearly seen that the overpotential for the parasitic reaction increases more quickly with the current density. This is the case for the competition between the chlorine evolution and water oxidation reactions. This simplified treatment assumes that the reactions take place on separate active sites, an assumption that might be questionable if the two reactions compete for the same active sites. It also neglects the effects of both mass transport of reactants to the anode surface, and of the formation of bubbles,247,248 on the local concentrations of reactants at the anode.247,248 Both of these effects will become important at the high current densities employed industrially. As several different reactions yield oxygen, especially under chlorate production conditions, and since the effect of competitive adsorption is important, this Tafel treatment of the influence of current density gives a limited understanding of the effect of current density. A more quantitative variant of this type of analysis is that by Tilak et al.,69 as given by eqs 26−28. Nevertheless, studies performed so far have found that increasing current densities decrease the selectivity for oxygen evolution (refs 54, 56, 57, 62, 66, 69, 70, 75, 162, 164, and 165). This could indirectly imply that the most important anodic sidereactions that yield oxygen have higher Tafel slopes than the chlorine evolution reaction. However, this effect is likely tangled with the effect of a decreasing pH at the anode surface with increasing current density. The effect of current density is probably that exhibited also by electrodes with decreased surface area,69,78,79 which yields an increase in local current density. Moreover, it is important to also recall that Eberil et al.74 found a maximum in chlorate formation efficiency close to the critical anode potential (for RuO4 formation), meaning that increased current density increases the selectivity only up to a certain point at chlorate conditions. Although several studies have seen effects of current density on the selectivity, further study of this factor under chlorate conditions is warranted. The effect of increasing concentration of hypochlorite species is related to the decomposition of hypochlorite species to form oxygen (refs 54, 55, 57, 71, 116, 121, 130, 162, and 170). This reaction will be discussed in more detail in section 3.2. An effect of chlorate concentration, due to competitive adsorption of chlorate on the electrode, on the oxygen evolution side reaction has been shown.54,162 However, Hammar and Wranglén54 made estimations that showed that the increase in electrolyte viscosity that chlorate brings could be the main reason for the decrease in oxygen evolution rate. This increase in viscosity would result in a decrease in the mass transfer rate of hypochlorite to the anode surface and a decrease in oxygen produced through decomposition of hypochlorite. The effect of temperature is probably complicated to elucidate in a detailed way, since it has effects on both thermodynamics and kinetics. This is true especially at chlorate cell conditions, since several reactions are involved in the main reactions yielding chlorate and the side reactions yielding oxygen. Several studies have found that increasing temperatures yield increased rates of oxygen evolution at chlorate conditions.54,57,162,164 This has been suggested to be related to increasing rate of diffusion of hypochlorite species to the anode surface at increasing temperatures54 and to increased rates of anodic and homogeneous decomposition of hypochlorite species at higher temperatures.162 Eberil et al.74 found a complicated temperature

(11)

It is obvious that the latter equilibrium is dependent on the pH of the solution. The reversible potential of the first reaction is also influenced by the chloride concentration, as will be discussed. Oxygen evolution is increasingly unfavorable as the pH decreases, keeping other conditions (e.g., current density) unchanged, due to the pH dependence of the electrode potential for oxygen evolution, unlike that for chlorine evolution.41,151,165 Only in chlor-alkali processes with HCl addition can oxygen evolving reactions be ruled out almost completely.41,151 This would imply that oxygen evolution is a more severe problem in chlorate processes, in which the pH of the process solution bulk is higher (pH 5.5−7). However, Trasatti and Lodi165 point out that if water oxidation occurs at the anode, the water oxidation reaction 11 acidifies the electrolyte close to the anode. Even if no water oxidation occurs, the formation of hypochlorous acid, reaction 4, acidifies the electrolyte. It is therefore considered reasonable that, even in processes where the bulk pH is high, the pH at the anode surface is in the acidic range. However, this reasoning is not sufficient to actually rule out any of the oxygenforming reactions, since the pH is probably not sufficiently acidic at the anode surface to fully prevent water oxidation. Indeed, several studies54,56,67,162,165 have shown clear correlations between pH and the oxygen evolution rate. The disagreement between different studies of the influence of pH also highlights an important difficulty in studying pH effects in chlorate electrolytes. Hammar and Wranglén54 did not find any influence of pH on the oxygen evolution rate between pH = 6.5 and pH = 11. However, Jaksić56 later found that the optimal pH for chlorate production should be in the pH = 6−6.5 range. Jaksić believed this disagreement to be the result of improper handling of the buffering effect of carbonates from absorption of CO2 from the air. Moreover, Hardee and Mitchell162 only found an influence of pH at pH > 7, and no effect of pH on the rate of oxygen evolution in the range of pH = 5−7. However, Hammar and Wranglén54 used graphite anodes. Nevertheless, chlorate electrolytes are buffered both by dichromate species and hypochlorite species, and several reactions in the process involve H+, making the study of pH effects in chlorate electrolytes somewhat difficult. Eq 2 also shows the influence of chloride concentration. Several studies and reports66,70,75,76,121,128,154,155,162,170,173 note that increasing concentrations of Cl− increase current efficiencies for chlorine or chlorate production and decrease the selectivity for oxygen evolution. However, this might not only be related to the decrease in the reversible electrode potential for chlorine evolution but also due to increased rates of mass transfer of chloride to the anode surface as well as competitive adsorption of chlorides on the surface of the anode, preventing anodic oxygenevolving reactions.66,165 The next fundamental factor that needs to be considered is that of the current density on the selectivity. It is simple to consider the effect of current density on two separate reactions, such as reactions 2 and 11. If Tafel expressions are written for both reactions, with the respective subscripts m and p for the main (ClER) and the parasitic (OER) reaction, the following expressions are obtained ηm = A m + bm × log(j)

(44)

(43)

for the main reaction and 3016


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evolution under chlorate process conditions, but they found that addition of nitrates resulted in increased oxygen evolution rates. The effects of all of these species seem to be related to competitive adsorption on the anode surface, and these effects might become less important at high chloride concentrations. Dichromate has been reported to both increase137 and decrease173 the selectivity for oxygen evolution. Interestingly enough, other studies have seen only minor effects.162 All three studies were performed at chlorate process conditions. Studies of dilute hypochlorous acid solutions have not found any effect of dichromate on the selectivity for oxygen evolution.92,95 A small number of metals, salts, and oxides of the elements have been shown to exhibit catalytic effects on hypochlorite decomposition. These are Co, Ni, Ir, and Cu. The effects of Fe, Mn, Hg, and Sn are less clear, with some (especially very old studies) reporting acceleration of hypochlorite decomposition. Connick and Hurley99 suggested that ruthenate ions might catalytically decompose hypochlorite ions (eqs 16 and 17). However, considering that the window of stability for the ruthenate ion is at high pH values of about 13−14 and potentials less than 1 V versus SHE,214 the prevalence of this reaction as an explanation for the oxygen evolution side reaction in either the chlor-alkali or the chlorate process is deemed quite low. No studies have been performed to compare the effects of different elements under industrial chlorate conditions. Some materials have not been studied since the 1800s, and a more general study of the heterogeneous activities of different metals, and possibly combinations of different metals, could serve to settle these issues. An important aspect to consider in such studies is the speciation of the metallic compounds, as it is possible that salts that are added will be oxidized to form metallic oxides active for decomposition of hypochlorite.92 It is clear that the influence of many different possible electrolyte contaminants and additives has not yet been examined fully under industrial conditions. 3.1.3. Influence of the Anode Structure and Composition. The effect of the anode properties on the oxygen evolution side reaction has been studied by many. Unfortunately, many studies lack important details regarding the properties of the anode. For example, proprietary coatings of unnamed composition that lead to lower oxygen evolution rates have been used in several experimental studies.36,162,167,173 Most studies have focused only on the selectivity between oxygen and chlorine formation under chlor-alkali conditions. An exception is the study of a modified coating with decreased selectivity for oxygen evolution at chlorate conditions performed by Spasojević et al.71−73 Additionally, detailed characterization, such as of the actual composition of the final coating has usually not been reported in literature, even if studies on proprietary anodes are excluded. This is a problem, especially since the preparation of anodes used for research usually is a manual process where the coating is, for example, brushed onto a substrate. Furthermore, it is known that the deposition efficiency of different compounds can vary widely depending on which precursors and which preparation conditions are used.249 This lack of detailed materials characterization could explain some of the conflicting results in literature. This does not necessarily imply that the studies are poorly done, since many of the studies were performed when characterization methods that are common today, such as SEM, XRD, XPS, and atomic force microscopy (AFM), had not entered general use. Furthermore, the usage of other X-ray spectroscopies (e.g., XAS), which allow a deeper understanding of factors such as the local structure, were not as accessible as today. Nevertheless, even studies conducted today often do not

dependence under chlorate conditions, at which the temperature had different effects at anode potentials on different sides of Ecr. Furthermore, the critical anode potential was associated with a maximum in current efficiency for chlorate production and a minimum in oxygen evolution current efficiency. This shows that it is imperative that further study of the selectivity issue is performed with equipment that allows precise correction of iR drop, since the critical anode potential is reached only at high applied current densities. Using such equipment makes it possible to determine the effect of Ecr clearly. The effects of anode geometry that Wanngård57 found is likely related both to a decrease in local current density, due to the increased geometric surface area of the higher electrode blades, and a change in the concentration gradient in the electrolyte flow across the blade for the higher type of blade (the “lower” blades were 0.333 m high and 0.379 m wide, while the higher blades were 0.8 m high and 0.3 m wide). Increasing the external holding volume has been found to increase the charge yield for chlorate formation.55,62 This is likely related to the increased residence time for the electrolyte afforded by a larger holding volume, enabling a higher conversion in the homogeneous chlorate formation reaction. 3.1.2. Influence of Additions and Contaminants in the Electrolyte. There have been comparatively few studies on the effect of contaminants and additions to the electrolyte on the oxygen evolution side reaction. In some cases, the species that have been studied have been found to have opposing effects. The great majority of studies are performed on systems that differ greatly from industrial chlorate or chlor-alkali conditions (e.g., on bleach or hypochlorous acid solutions), again indicating the importance of performing further studies under conditions similar to those used industrially. Fluoride was found by both Fukuda et al.144 and Baolian and Wenhua173 (based on the overall chlorate cell current efficiency) to decrease the rate of oxygen evolution. On the other hand, Jaksić et al.137 found essentially no influence of fluoride on the oxygen evolution rate under chlorate conditions. The difference between the study of Fukuda et al. and Jaksić et al. might be possible to rationalize based on the differences between the electrolytes studied (Jaksić et al. studied a chlorate electrolyte while Fukuda et al. studied an acidic electrolyte), if the results of Baolian and Wenhua were ignored. Still, the results might be reconciled by noting that the effect seen by Baolian and Wenhua was significantly smaller (about 1%) at 200 g/dm3 NaCl than at 50 g/dm3 NaCl. Furthermore, the measurements of Jaksić et al. were performed with a NaCl concentration of 310 g/dm3. However, Wanngård168 notes that F− needs to be avoided since it also has the effect of deactivating the anode. Kelsall and Robbins145 have predicted, based on thermodynamic calculations, that F− addition should result in dissolution of TiO2 to − form TiF2− 6 . Still, it might be concluded that the effect of F addition on selectivity essentially is that of competitive adsorption on the electrode surface, reducing the rate of oxygen formation in dilute electrolytes, but also possibly competing with the adsorption of Cl−. Phosphates have been found to increase the rate of oxygen evolution.137,157,173 In contrast, the PdSn2-containing coating devised by Spasojević et al. made use of the adsorption of phosphates on Pd to prevent oxygen evolution. Sulfates have been reported to decrease the efficiency of chlorine evolution.157,165 Kazarinov and Andreev143 concluded that both phosphates and sulfates adsorb strongly on RTO, displacing chlorides. However, Jaksić et al.137 saw no effect on oxygen 3017


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DFT studies by Rossmeisl et al.,251 Hansen et al.,183 and of Exner et al.207,210 In these studies, the importance of adjacent CUSs on the anode surface were considered. First, it has been found that it is not only the metal CUSs themselves (found by Hansen et al.183 to be the active sites on weaker-binding oxides) but also Ocovered CUSs183,207 that are active sites for both ClER and OER. Second, it has been found that the mechanism for OER does not necessarily involve O−O recombination,183,207 thus not being dependent on paired CUSs. The finding that the adsorption energy of O on CUSs on the surface is the deciding factor for both chlorine and oxygen evolution activity might at first be hard to reconcile with experimental studies that show that materials with increased overpotential for oxygen evolution also show an increase in selectivity for chlorine evolution. This can be understood by studying Figure 23, which shows that there are ΔE(Oc) values which allow decreased selectivity for oxygen evolution, compared with RuO2, with maintained low overpotential for chlorine evolution.64 Indeed, Karlsson et al.64 have built upon the results of Hansen et al.183 to conclude that the increased selectivity exhibited by RTO electrodes likely is caused by an electronic activation of Ti by Ru. The effect on selectivity from changing the relative amounts of Ru and Ti in the binary coating is thus not a matter of separation of active Ru sites67,68 but related to the fundamental electronic properties of the mixed oxide coating and its effects on the activity of Ti sites. Much work has been devoted to improving the selectivity by using other materials than RTO, or by doping RTO. Studies which examine the effect of IrO2 have found it to have a lower selectivity for oxygen evolution than pure RuO2,70 and Ir-doped RTO have also shown decreased O2 selectivities.74,165 At the same time, the dissolution of Ir compounds from the anode might yield an increase in oxygen evolution in the industrial setting due to catalysis of homogeneous hypochlorite decomposition. However, RuO2−Sb2O5−SnO2 electrodes have been found to exhibit similar selectivities in dilute solutions regardless of Ir doping.185,186 PdO has been reported to be a stable and highly selective ClER anode.148 Shlyapnikov141 found that Pt and RuO2 exhibited higher current efficiencies for chlorate formation than graphite, lead oxide, or manganese oxide anodes. On the other hand, platinized Ti has been reported to have a lower chlorine evolution efficiency than RuO2−SnO2-coated Ti.187 Co3O4 and related compounds have also been reported to have a lower selectivity for oxygen evolution.78,140,147,165 However, also in this case, Co corroded from the surface of the electrode might yield increased oxygen evolution due to homogeneous hypochlorite decomposition. Combining Ru with Sn in oxide electrodes has been found to have generally similar effects as the combination of Ru and Ti in RTO electrodes.172 However, studies have found that Ru−Sn oxide electrodes are less selective for OER than RTO electrodes.69,188 Tilak et al.69 noted similar effects when combining Ru and Hf in mixed oxide electrodes. Doping RuO2 with Zn, Ni, or Mg have also been found to yield lower selectivities for oxygen evolution, in comparison with those of pure RuO2.77,181,182,193 The insights from DFT63−65,183,184,207,210 are also interesting when discussing the recent study of Spasojević et al.60 They found that Pt-IrO2 overlayer coatings on RTO show decreased selectivity for OER as a side reaction in both chlor-alkali and chlorate electrolytes.60 The overlayer electrodes exhibited a significantly lower electrochemically active surface area (increased local current densities), which is likely part of the reason for the improvement in selectivity (see section 3.1.1). However, the work might also be reflective of the same fundamental effects

make use of the possibilities afforded by modern characterization methods. All the same, some overall factors can be identified. One effect, that of the active surface area of the electrode, is likely primarily connected to its effect on the local current density, as discussed in section 3.1.1. While a large active surface area is obviously important to achieve a high utilization of the coating, an increased electrode surface area is also associated with a decreased local current density and resulting increased selectivity for oxygen evolution. An increased electrode active surface area is often achieved by decreasing the size of the particles making up the coating. The selectivity effects from doping RuO2 or RTO that are found in many studies are most likely due to this effect. Arikawa et al.75 found that the selectivity is affected by the RTO loading on the anode, with increasing loading resulting in increased selectivity for OER. This effect is possibly also connected to an increase in active surface area, especially since Arikawa et al.75 studied electrodes with comparatively low loading. The atomic-level electronic effect of using different combinations of cations in mixed oxide electrodes (e.g., the combination of Ru and Ti in Beer-type DSA) is more complex. Certain trends that are valid for Ru-based oxides might not apply for, for example, Ir-based oxides, as has been shown by, for example, Kuznetsova et al.189 (pure IrO2 being more Cl2-selective than any of the doped oxides considered, while Ti or Sn-doped RuO2 oxides are more Cl2-selective than pure RuO2). Even if one focuses only on “Ru-based” electrodes, the electronic structure and therefore also the electrochemical properties of materials containing low concentrations of RuIV are significantly different from materials containing high RuIV concentrations.64 However, it is important to consider the cautionary observation of Trasatti156 that the selectivity of different anode materials does not generally exhibit varying selectivities for chlorine evolution versus oxygen evolution. Materials highly active for chlorine evolution are usually also highly active for oxygen evolution. A first fundamental understanding for this behavior has been provided by the DFT studies of Hansen et al.,183 showing that ΔE(Oc) determines the activity for both chlorine and oxygen evolution reactions on several metallic oxide materials. The fundamental factor that governs the selectivity is the descriptor value placing each material in a certain part of the activity volcano relevant for a certain set of linearly scaling reactions.183,199,250 Nevertheless, while the same factor governs both reactions, there is some room for optimizing the selectivity for chlorine evolution over oxygen evolution even on conventional rutile oxide electrodes, as has been explored theoretically63−65 and as experimental studies also indicate (see Table 6). The effect of altering the relative amounts of Ru and Ti in RTO coatings will be discussed first. A decrease in the Ru content has consistently been shown to result in decreased selectivity for oxygen evolution (refs 67, 68, 75, 159, 161, 165, 167, and 171). Similar trends regarding Ru content and OER selectivity have been noted for Ru−Sn oxide electrodes.172 This might be another aspect on the observation that an increased overpotential for oxygen evolution of an anode material (i.e., a decreased j0 for reaction 11) has been linked to a decreased selectivity for oxygen evolution in chloride-containing solutions.69,150,159,167 These factors have been related to mechanistic differences between the oxygen-evolving reaction and the chlorine-evolving reaction. It has been proposed that both reactions need Ru sites to progress, but that the oxygen-evolving reaction needs two Ru adjacent sites.67,68 Here an interesting new insight can be gained from the 3018


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discussed in several theoretical works.63−65,183,210 Doping TiO2 with Ru changes the descriptor value of the doped oxide toward the part of the volcano optimal for chlorine evolution,64 and it is possible that Spasojević et al.60 is noting a similar effect when mixing IrO2 with Pt. PtO2 and IrO2 are located on opposite sides of the volcano tip optimal for selective and active chlorine evolution.65,183,210 Some have found effects on the selectivity for chlorine evolution based of the support material.179 This influence is also hinted at in the original Beer patents.8,22 Less attention has been devoted to the possibility that the electrode might heterogeneously catalyze the decomposition of hypochlorous acid species. While certain compounds (such as Co) are generally not considered as electrode components, due to their well-known catalytic effect as electrolyte contaminants (see the previous section), there are fewer studies of heterogeneous hypochlorite decomposition. When it comes to metallic and oxidic materials (see Table 5), Co, Ni, Ir, and Cu are materials which have been found to catalyze the bulk decomposition. Still, Kotowski and Busse159 did not find that Pt/Ir particles were active for hypochlorite decomposition, despite the Ir content. Kuhn and Mortimer136 did see indications that RuO2 heterogeneously decomposes hypochlorite in a Cl2saturated solution, while Kokoulina and Bunakova155 and Kotowski and Busse159 found no effect when RTO electrodes were used. To get complete certainty in this matter, studies where both RuO2, RTO and doped RTO or RuO2 electrodes are tested under the same conditions, with the same equipment and methods, are likely required. Finally, some studies36,74,141,150,154 indicate that for the chloralkali process, it might not be enough to simply optimize anodes for minimal oxygen selectivity, as this is then associated with increased rates of chlorate production. This might partly be a pH effect, as the OER decreases pH at the anode and therefore also decreases the rate of chlorate formation.154 A reduced rate of the OER with a low-oxygen anode might thus require additional anolyte acidification,36 although this might yield an increase in cell voltage due to decreased conductivity of the cationpermeable membrane.2 Whether the effect could be associated with changes in activity or selectivity for electrochemical chlorate formation on the surface of the anode is yet to be determined.

Finally, Tilak and Chen5 believe it has an insignificant influence under chlorate production conditions. The influence of the water oxidation reaction, reaction 11, is also uncertain. It is known that this reaction dominates at low chloride concentrations and low electrode potentials,54 but at actual industrial current densities and concentrations, the influence is unclear. Kotowski and Busse159 think it makes up a significant amount of the oxygen evolved in the membrane chloralkali process, while Hardee and Mitchell162 and Spasojević et al.71,72 find it to be insignificant at chlorate conditions. Tilak and Chen5 believe it is the only important source of oxygen at chlorate conditions. The studies of Spasojević et al.71 and Spasojević et al.60 indicate that a majority of the oxygen is formed through water oxidation (see Figures 8 and 18) in dilute solutions with similar hypochlorite concentrations as in chlorate production. It seems that the rate of water oxidation is significantly lower in concentrated chlorate solutions with chromate buffering,57,71,72,162 making it a less important source of oxygen than the decomposition of hypochlorite in sodium chlorate production. The water oxidation reaction is likely kinetically controlled, while the oxidation of hypochlorite species is diffusion controlled,60,71,72 at least at the current densities applied in industrial chlor-alkali and chlorate production. The homogeneous decomposition of hypochlorous acid to form oxygen gas, reaction 14, is well-known. The catalytic effect of different electrolyte contaminants on this reaction has been discussed in section 3.1.2. Hypochlorite does decompose to form oxygen even in the absence of any catalyst, although the rate is significantly lower than that of chlorate production.89,91,92 There are indications that both reactions share an intermediate.89,92 While measures to control the electrolyte pH can be applied to prevent the formation of hypochlorous acid species in the chloralkali process,36 this is obviously not a possible recourse in the chlorate process. Further research is needed to elucidate the connection between oxygen and chlorate formation, so that measures might be found to influence the selectivity in the direction of chlorate formation. More studies of this reaction are needed to fully understand its importance in chlorate production. Still, it can be said that purification of the electrolyte to remove metallic contaminants should minimize the importance of homogeneous hypochlorite decomposition in both chlor-alkali and sodium chlorate production. Finally, the influence of anodic decomposition of hypochlorite species, resulting in formation of oxygen but not chlorate, is controversial. Kotowski and Busse159 have suggested one mechanism for this reaction, eq 23. Older studies consider the anodic decomposition of hypochlorite a possible source of oxygen,116,121,130 but most studies performed before 1980 seem to conclude that it is a minor source of oxygen. Several studies performed since the end of the 1980s have shown it to be an important factor for oxygen evolution in chlorate systems.54,57,159,162,177 The fact that the anodic decomposition of hypochlorite was not thought to be an important side reaction at chlorate process conditions until the second half of the 1980s shows that indirect methods used to measure current efficiency for chlorate production are insufficient to give a full understanding of the process. It took studies where the proper detection of the oxygen evolved was performed to find that other reactions might also be important sources of oxygen. Nevertheless, some still do not agree that anodic decomposition of hypochlorite species is important,5,167 warranting further studies of how hypochlorite is actually decomposed at the anode. The relative importance of heterogeneous chemical decomposition

3.2. Extent of Oxygen Evolution from Different Reactions

We now change focus to discuss the relative importance of the different reactions thought to result in oxygen production. Only a few studies have attempted to quantify the five most probable oxygen-evolving reactions, reactions 10, 11, 12, 14, and 23, that are present in varying degrees at chlorate and chlor-alkali conditions. Due to the low pH at the anode surface, anodic hydroxide oxidation, reaction 12, is generally assumed to be an unimportant source of oxygen. The importance of the other reactions is less clear. The influence of the first reaction, the anodic chlorate formation reaction, is controversial. Studying chlorate process conditions, Hammar and Wranglén,54 Ibl and Vogt,62 and Jaksić et al.55 thought it was the most important oxygen-evolving reaction. Recent modeling studies make the same assumption.191 These models describe the overall current efficiency of the process well. However, Hardee and Mitchell162 seem to think it has a small influence. Studying a membrane cell chlor-alkali system, Kotowski and Busse159 thought it could make up anywhere from 20% to 60% of the oxygen evolved in the system. 3019


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at both acidic, neutral, and alkaline pH values, while the solubility of pure TiO2 is predicted to be low under such conditions.145 Despite these predictions for the pure oxides, the lifetime of industrial RuO2−TiO2 mixed-oxide DSAs is about 10 years.253 Additionally, the anode potential of electrodes in operation does not change until the RuIV content has become very low.226 The constancy in activity during corrosion could perhaps be seen as support for our recent suggestion that also Ti sites can be active in DSAs and that only low concentrations of Ru are necessary to achieve these effects.64 Perhaps it might even be conceivable that the last sites that are active on DSA are Ti sites activated by Ru, as these sites might be more stable.65 While not directly relevant to the selectivity issue, some conclusions can be mentioned regarding the supposed deactivation of DSA due to formation of a nonconductive interlayer of TiO2, between the support and the coating, during electrolysis.223,236,237,239,254 This deactivation mechanism is frequently mentioned, but there seems to be little direct proof of it being an important deactivation mechanism for electrodes used in chlor-alkali or sodium chlorate production. It has been suggested that the deactivation is more likely the result of complete removal of Ru from the surface layers of the coating.226,227 Indeed, the fact that reactivation by Pt deposition onto the surface of used industrial electrodes seems to be possible supports this claim.2 We would like to point out that much of the detailed knowledge about the connection between selectivity and stability has resulted from the radiochemical studies performed almost exclusively by scientists in the Soviet Union between 1960 and 1990. These methods have allowed the corrosion of the anode to be studied in situ during electrolysis at industrially relevant current densities and electrolyte composition and not only based, for example, on an increase in potential in accelerated lifetime tests. The latter type of stability test, which gives very little understanding of the character of the degradation during electrolysis, has been the dominating one employed in the West even up until year 2015. The reason is perhaps the practical challenge in employing the radiochemical method, as proper safeguards must be employed, and possibly also the long time periods needed to study anode degradation at industrially relevant conditions. The study of Gajić-Krstajić et al.,239 which employed a spectrophotometric method to allow similarly sensitive detection of Ru as afforded by the radiochemical method, could perhaps be used to avoid handling radioactive materials and give time-resolved understanding of the degradation process. Furthermore, the recent studies of Mayrhofer and co-workers240−243 and Chorkendorff and co-workers245,246 have shown how ICP-MS and EQCM can be used to monitor degradation of DSAs during use as OER catalysts. Nevertheless, long-term tests in the correct conditions, so that the effect of the relevant reactions are studied, are still necessary. Additionally, the clear link between OER and degradation that the literature indicates means that, for example, an accelerated lifetime test in sulfuric acid, where OER is the only reaction occurring, constitutes a test of a completely different property than what is relevant for an anode which is to evolve almost exclusively chlorine gas. As seen in several studies (refs 74, 164, 174, 219, 222, 223, 231, and 238), there is a change in the Tafel slope for OER and ClER on Ru-based anodes at the critical anode potential, Ecr. This potential is still not well understood but has been related to the formation of RuO4.219 Indeed, the Pourbaix diagram indicates the connection between Ecr and the start of formation of

and electrocatalytic decomposition of hypochlorite species on, for example, the RTO electrode surface, is also an open question. Furthermore, in studies where the influence of anodic hypochlorite decomposition has been considered under chlorate-process-like conditions, some studies find that the rate of oxygen evolution shows characteristics of mass transport limitations (even on graphite),54 while others find it shows characteristics of kinetic limitations (on RTO).57 Studies have found that the reaction could make up as much as 40−80% of the overall oxygen production, even in membrane chlor-alkali cells.159 Wanngård57 found that if the homogeneous decomposition reaction can be avoided (reaction 14), the anodic decomposition of hypochlorite (the reaction where hypochlorite is decomposed electrochemically on the anode surface, proposed by Kotowski and Busse159 to proceed via reaction 23) is the main contributor to the oxygen evolved under chlorate conditions.57 However, the actual mechanism of the decomposition, whether hypochlorous acid or hypochlorite is the main reactant, is still not known. Physics-based macroscopic modeling has suggested that hypochlorite (OCl−) is the most probable reactant,176 while estimations based on a probable low pH at the anode surface suggest that HOCl should be the reactant.170 The recent study of Macounová et al.,84 suggesting the possibility of radical reactions involving ClO· during oxidation of hypochlorite, also indicates the need for further research on the topic. 3.3. Connection between Anode Stability and Selectivity

The literature clearly shows the connection between the rate of DSA degradation and the rate of the OER. However, it must be pointed out that the relationship is more complex than it may seem. In chloride-containing electrolyte, at very low pH values below 2, degradation due to the formation of soluble chlorides is likely a dominating mechanism of anode degradation.213 At these pH values, the corrosion rate of Ti has been found to exceed that of Ru.229 The rate of corrosion of DSA has been found to be minimal in the pH range of 2−3,151 which agrees with the area of maximal stability according to the Pourbaix diagram.214 As the pH is increased, the rate of corrosion increases in a similar rate as the OER.163,217,218 In this area, the corrosion rate of Ti has been found to be constant, but the Ru corrosion rate increases significantly.229 It may be noted that this increase coincides with the chlorine-hypochlorite equilibrium being pushed toward hypochlorite, see reaction 4 and Figure 1. The oxidizing power of hypochlorite, and the fact that hypochlorite, as opposed to chlorine, does not form gas bubbles in the electrolyte, may accelerate the oxidation of ruthenium. There is a clear link between OER and RuIV corrosion in the pH range relevant to especially industrial chlor-alkali production (pH 2 to 4). While less data is available under pH values suitable for chlorate production, it is likely that the trend is maintained even there.231 However, there is no clear link between the rate of ClER and that of RuIV corrosion in the pH range relevant for chlor-alkali or chlorate production. At higher pH values (above 10), the rate of corrosion of both TiIV and RuIV increase,229 possibly the result of formation of metal oxyhydroxides.142,213,230 Literature indicates that the surface layers of electrodes used in chlor-alkali or sodium chlorate production will be stripped of RuIV. This has been shown directly both radiochemically229 and by, for example, selective ion mass spectrometry or AES,235−237,252 and indirectly by other means.226,227 This is in agreement with thermodynamic predictions,214 which indicate that soluble Ru compounds should form during anodic polarization of pure RuO2 in concentrated chloride solutions, 3020


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ruthenate species.214 Furthermore, an increase in the rate of Ru dissolution from DSA has been noted at electrode potentials above Ecr.231 Thus, operation at anode potentials higher than Ecr should be avoided to limit DSA corrosion. From an industrial perspective, the risk of exceeding Ecr is higher in the chlorate process than in the chlor-alkali process.238 This is partly connected to the decrease in Ecr which occurs as pH is increased223 and which in turn likely is related to the pH effect on the potential required for ruthenate formation, which follows the same trend.214 Interestingly, Eberil et al.74 found an optimum in selectivity for chlorate production close to Ecr. Another important industrial aspect is the finding that Ru corrosion increases by repeated interruptions in the electrolysis.224 Thus, repeated starts and stops should be avoided. In closing, we would like to provide a discussion of how the lifetime of industrial DSA can be maximized in the intermediate pH range relevant for chlor-alkali and sodium chlorate formation. The difference in corrosion behavior between the Ru and Ti component of DSA in the intermediate pH range is most likely due to the possibility of forming ruthenates and perruthenates.213 Ti does not form such compounds.230 The formation of gaseous RuVIII products during electrolysis using DSA is well-known both in literature (see, for example, refs 151, 154, 221, 224, and 229) and in industrial chlor-alkali and sodium chlorate production. This aspect has devoted significant experimental work, with, for example, the study of Kötz et al.212 clearly showing the connection between the onset of OER and the onset of RuO4 formation on pure RuO2 electrodes. It is reasonable to assume that the formation of RuO4 occurs via an intermediate that is shared with the OER, as surface oxides need to form in both cases. Both the rate of OER and the percentage of the Ru released into the gas phase increases with pH.151,154 Furthermore, it was clearly shown by Pecherskii et al.67 that Ru electrodes that exhibit improved selectivity for ClER also exhibit improved stability. Even for other Ru mixed oxides, such as Ru−Sn oxides, improved chlorine evolution selectivity69,71,172 is associated with improved stability.234 RuIV corrosion rates have been found to decrease significantly with increasing Cl− concentration.154,218,231 As has been discussed in the preceding sections, increasing the Cl− concentration is a very simple way of increasing the current efficiency for Cl2. Increasing the current density is another way of improving the current efficiency for Cl2, at least as long as the density is not so high as to, for example, result in mass transport limitations, and this also results in decreased current efficiency for corrosion.154,221,224,231 Increasing the temperature decreases the selectivity for chlorine evolution and increases the rate of corrosion of DSA.154,163 However, a higher temperature might still be viable, as it improves the kinetics of the chlorine- and chlorate-producing reactions and decreases the cell potential difference. At pH values relevant for industrial chlor-alkali and sodium chlorate formation, high oxygen production rates are indicative of a high rate of RuIV corrosion due to the formation of ruthenates or perruthenates. Thus, chlor-alkali production should operate in the pH range of 2−3 to achieve a maximal chlorine current efficiency41,151 and anode lifetime.151 For chlorate production, the optimal pH for current efficiency of chlorate production is significantly higher, at about pH 6−7. In this pH range, it is not possible to maximize both chlorate production efficiency and DSA stability in the same way, and both factors are connected to the critical potential. Anodes in both processes should be operated below the critical potential.

4. CONCLUSIONS It is clear that the issue of the selectivity between the main reactions in chlor-alkali and chlorate processes and the oxygenevolving side reactions requires further study, as has also been recommended by Wanngård168 and Tilak and Chen.5 While the effect of several factors, most importantly of the current density, are relatively well-known, the details of the effects of process parameters, electrolyte composition, and electrode characteristics on the selectivity issue, as well as the mechanisms and relative importance of the suggested oxygen-forming reactions, are still not well-elucidated. Several studies have been conducted under conditions similar to those used in industrial chlor-alkali production, while fewer studies have examined the selectivity under conditions similar to those used during industrial chlorate production. For chlor-alkali production, a practical way of increasing the selectivity through acidification of the electrolyte is available for state-of-the-art membrane cells. However, a higher pH is necessary to facilitate chlorate formation, thus excluding the method of acidification in industrial chlorate production, and oxygen evolution still accounts for about 5% current efficiency loss. Additionally, the possible influence of the critical anode potential, Ecr, on the selectivity for chlorate production is still not well understood. Its appearance at high, industrially relevant, current densities makes it imperative that future studies employ accurate iR compensation methods. Such methods could then be combined with measurements of both the gas and the liquid phase composition which might also give further understanding for the relative importance of the different oxygen-evolving reactions. For both processes, modern experimental and theoretical methods should be used to gain a deeper understanding of the interplay between anode activity and selectivity, composition, and structure. The combination of theoretical modeling and modern characterization methods (e.g., the different X-ray spectroscopic methods available today) is well-suited to exploring the details of heterogeneous and homogeneous (electro)catalytic reactions. These tools now allow for detailed understanding of the electronic structure of practical catalysts, whereas before mostly macroscopic parameters regarding structure and composition of electrodes have been possible to study in detail. The understanding resulting from studies combining theory and experiments could then be used to improve the selectivity and activity of electrodes used in chloralkali and sodium chlorate production. They could also be used to start exploring the details of the relatively poorly understood bulk-phase reactions, where catalytic processes are involved both in unwanted oxygen evolution reactions as well as in the formation of sodium chlorate. The latter aspect is especially interesting, as an understanding of the catalysis of sodium chlorate formation could result in new ways of accelerating the reaction, with important consequences for sodium chlorate process design. AUTHOR INFORMATION Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest. Biographies Rasmus K. B. Karlsson received his Master’s degree in Chemistry and Chemical Engineering in 2011 and his Ph.D. in Chemical Engineering in 3021


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CV

2015, both from KTH Royal Institute of Technology. His Ph.D. thesis research was primarily focused on combining experimental and theoretical methods to elucidate the fundamental mechanisms deciding the selectivity between chlorine and oxygen evolution in the chlor-alkali and chlorate processes.

DEMS DSA EDX EIS EQCM EXAFS

Ann Cornell received her MSc in Chemical Engineering in 1987 and then worked for 15 years in industry with research and development within electrochemical engineering, mainly related to industrial chlorate electrolysis. She completed her Ph.D. in 2002 at KTH with a thesis on electrode reactions in the chlorate process and has, since 2003, worked in the group of Applied Electrochemistry at KTH. Ann is now an associate professor and still conducts research related to industrial electrolysis but has also broadened her interests to areas such as cellulose-based Li-ion batteries.

GC ICP-MS MS OER ORTA

ACKNOWLEDGMENTS The financial support of the Swedish Energy Agency and of Permascand AB are hereby acknowledged. Furthermore, the constructive comments given by Dr. John Gustavsson (Permascand AB) on a draft of this review are gratefully acknowledged. Staffan Sandin is acknowledged for bringing some of the papers on decomposition of hypochlorous acid species to our attention. Professor Kiyotaka Asakura (the Catalysis Research Center of Hokkaido University, Japan) is acknowledged for assistance with providing a high-resolution version of Figure 8. Professor Lars G. M. Pettersson (Stockholm University) is gratefully acknowledged for constructive comments on section 2.2.5.2.

RDE RHE

RTO SCE SECM SEM SHE SIMS TEM TNT XANES XAS XPS XRD

NOMENCLATURE η overpotential, V Φep current efficiency (or charge yield) of product p ac adsorbate a on a CUS of the rutile (110) surface abr adsorbate a on a bridge site of the rutile (110) surface ai activity of ion i ci concentration of species i, mol/dm3 E electrode potential, V versus SHE, or energy (e.g., of adsorption), kJ E0,i equilibrium (or reversible) electrode potential for reaction i, V versus SHE Ecr critical anode potential G Gibbs free energy, kJ/mol iR compensation Correction for the electrolyte resistance in electrochemical measurements j current density, A/m2 ji partial current density for a certain reaction, A/m2 j0 exchange current density, A/m2 jcr critical current density, the current density at Ecr q* electrochemically active surface area, C/m2 t temperature in degrees centigrade U electrode potential AES atomic emission spectroscopy AFM atomic force microscopy BEP Brønsted−Evans−Polanyi CHE computational hydrogen electrode ClER chlorine evolution reaction CUS coordinatively unsaturated site

cyclic voltammetry differential electrochemical mass spectroscopy dimensionally stable anode energy-dispersive X-ray spectroscopy electrochemical impedance spectroscopy electrochemical quartz crystal microbalance extended X-ray absorption fine-structure spectroscopy gas chromatography inductively coupled plasma mass spectrometry mass spectrometry oxygen evolution reaction ruthenium−titanium oxide anode, abbreviation frequently used in Soviet and Russian literature rotating disc electrode relative hydrogen electrode, the SHE value adjusted for the pH of the solution (i.e., that measured with a SHE immersed in the same solution as the working electrode) ruthenium−titanium oxide saturated calomel electrode scanning electrochemical microscopy scanning electron microscopy standard hydrogen electrode secondary ion mass spectrometry transmission electron microscopy TiO2 nanotube X-ray absorption near-edge spectroscopy X-ray absorption spectroscopy X-ray photoelectron spectroscopy X-ray diffraction spectroscopy

REFERENCES (1) Schmittinger, P.; Florkiewicz, T.; Curlin, L. C.; Lüke, B.; Scannell, R.; Navin, T.; Zelfel, E.; Bartsch, R. Ullmann’s Encyclopedia of Industrial Chemistry; Wiley-VCH Verlag GmbH & Co. KGaA: Weinheim, Germany, 2012. (2) O’Brien, T. F.; Bommaraju, T. V.; Hine, F. Vol. I: Fundamentals; Springer, 2005; Vol. 1. (3) Schlag, S. Sodium Chlorate. Chemical Economics Handbook Product Review; 2012. (4) Fauvarque, J. The Chlorine Industry. Pure Appl. Chem. 1996, 68, 1713−1720. (5) Tilak, B. V.; Chen, C.-P. In Chlor-alkali and chlorate technology: R.B. MacMullin Memorial Symposium; Burney, H. S., Furuya, N., Hine, F., Ota, K.-I., Eds.; The Electrochemical Society, 1999. (6) Viswanathan, K.; Tilak, B. V. Chemical, Electrochemical, and Technological Aspects of Sodium Chlorate Manufacture. J. Electrochem. Soc. 1984, 131, 1551−1559. (7) Vogt, H.; Balej, J.; Bennett, J. E.; Wintzer, P.; Sheikh, S. A.; Gallone, P.; Vasudevan, S.; Pelin, K. Ullmann’s Encyclopedia of Industrial Chemistry; Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim, Germany, 2010. (8) Beer, H. B. Method of Making an Electrode Having a Coating Containing a Platinum Metal Oxide Thereon. U.S. Patent 4052271, 1977. (9) Roginskaya, Y. E.; Galyamov, B. S.; Lebedev, V. M.; Belova, I. D.; Venevtsev, Y. N. Phase Composition and Electrical Transport Properties of the RuO2-TiO2 System. Zh. Neorg. Khim. 1977, 22, 499−504. 3022


Chemical Reviews

Review

(10) Hine, F.; Yasuda, M.; Yoshida, T. Studies on the Oxide-Coated Metal Anodes for Chlor-Alkali Cells. J. Electrochem. Soc. 1977, 124, 500− 505. (11) Gerrard, W. A.; Steele, B. C. H. Microstructural Investigations on Mixed RuO2−TiO2 Coatings. J. Appl. Electrochem. 1978, 8, 417−425. (12) Over, H. Surface Chemistry of Ruthenium Dioxide in Heterogeneous Catalysis and Electrocatalysis: from Fundamental to Applied Research. Chem. Rev. 2012, 112, 3356−3426. (13) Kavan, L.; Grätzel, M.; Gilbert, S. E.; Klemenz, C.; Scheel, H. J. Electrochemical and Photoelectrochemical Investigation of SingleCrystal Anatase. J. Am. Chem. Soc. 1996, 118, 6716−6723. (14) Glassford, K. M.; Chelikowsky, J. R. Electronic Structure of TiO2:Ru. Phys. Rev. B: Condens. Matter Mater. Phys. 1993, 47, 12550− 12553. (15) García-Mota, M.; Vojvodic, A.; Abild-Pedersen, F.; Nørskov, J. K. Electronic Origin of the Surface Reactivity of Transition Metal Doped TiO2(110). J. Phys. Chem. C 2013, 117, 460−465. (16) Beer, B. H. Noble Metal Coated Titanium Electrode and Method of Making and Using It. U.S. Patent 3096272, 1963. (17) Beer, H. B. Electrode and Method of Making Same. U.S. Patent 3234110, 1966. (18) Beer, H. B. Electrolysis with Precious Metal-coated Titanium Anode. U.S. Patent 3236756, 1966. (19) Beer, H. B. Electrode and Coating Therefore. U.S. Patent 3632498, 1972. (20) Beer, H. B. Electrode Having Platinum Metal Oxide Coating Thereon, and Method of Use Thereof. U.S. Patent 3711385, 1973. (21) Beer, H. B. Method of Making An Electrode Having A Coating containing a platinum metal oxide thereon. U.S. Patent 3864163, 1975. (22) Beer, H. B. Coating of Protected Electrocatalytic Material on an Electrode. U.S. Patent 3933616, 1976. (23) Hayfield, P. C. S. Development of the Noble Metal/oxide Coated Titanium Electrode. Part I: the Beginning of the Story. Platinum Met. Rev. 1998, 42, 27−33. (24) Trasatti, S. Electrocatalysis: Understanding the Success of DSA. Electrochim. Acta 2000, 45, 2377−2385. (25) Duby, P. The History of Progress in Dimensionally Stable Anodes. JOM 1993, 45, 41−43. (26) Trasatti, S. In The Electrochemistry of Novel Materials; Lipkowski, J., Ross, P. N., Eds.; VCH Publishers, 1994; Chapter 5, pp 207−295. (27) Trasatti, S. Physical Electrochemistry of Ceramic Oxides. Electrochim. Acta 1991, 36, 225−241. (28) Daghetti, A.; Lodi, G.; Trasatti, S. Interfacial Properties of Oxides Used as Anodes in the Electrochemical Technology. Mater. Chem. Phys. 1983, 8, 1−90. (29) Electrodes of Conductive Metallic Oxides. Part B; Trasatti, S., Ed.; Elsevier, 1981. (30) Electrodes of Conductive Metallic Oxides. Part A; Trasatti, S., Ed.; Elsevier, 1980. (31) Trasatti, S. Progress in the Understanding of the Mechanism of Chlorine Evolution at Oxide Electrodes. Electrochim. Acta 1987, 32, 369−382. (32) Hayfield, P. C. S. Development of the Noble Metal/oxide Coated Titanium Electrode. Part II: the Move from Platinum/iridium to Ruthenium Oxide Electrocatalysts. Platinum Met. Rev. 1998, 42, 46−55. (33) Hayfield, P. C. S. Development of the Noble Metal/oxide Coated Titanium Electrode. Part III: Coated Titanium Anodes inWidely Ranging Oxygen Evolving Situations. Platinum Met. Rev. 1998, 42, 116− 122. (34) Over, H. Atomic Scale Insights into Electrochemical Versus Gas Phase Oxidation of HCl Over RuO2-based Catalysts: a Comparative Review. Electrochim. Acta 2013, 93, 314−333. (35) Diebold, U. The Surface Science of Titanium Dioxide. Surf. Sci. Rep. 2003, 48, 53−229. (36) Bergner, D.; Hartmann, M. Chlorate and Oxygen Formation in Alkali Chloride Membrane Electrolysis. J. Appl. Electrochem. 1993, 23, 103−107. (37) Standard Potentials in Aqueous Solutions; Bard, A. J., Parsons, R., Jordan, J., Eds.; CRC Press, 1985.

(38) Brinkmann, T.; Santonja, G. G.; Schorcht, F.; Roudier, S.; Sanco, L. D. Best Available Techniques (BAT) Reference Document for the Production of Chlor-alkali; JRC Science and policy reports EUR 26844 EN; Industrial Emissions Directive 2010/75/EU (Integrated pollution prevention and control), 2014. (39) Eurochlor, Chlorine Industry Review 2013−2014; 2014. (40) Wendt, H.; Vogt, H.; Kreysa, G.; Kolb, D. M.; Engelmann, G. E.; Ziegler, J. C.; Goldacker, H.; Jüttner, K.; Galla, U.; Schmieder, H. e. Ullmanna’s Encyclopedia of Industrial Chemistry [Online]; Wiley-VCH: Weinheim, Germany, 2009; Chapter Electrochemistry (accessed May 20, 2011). (41) Bergner, D. Reduction of By-product Formation in Alkali Chloride Membrane Electrolysis. J. Appl. Electrochem. 1990, 20, 716− 722. (42) Valensi, G.; Deltombe, E.; De Zoubov, N.; Vanleugenhaghe, C.; Pourbaix, M. In Atlas of Electrochemical Equilibria In Aqueous Solutions; Pourbaix, M., Ed.; Pergamon Press Ltd., 1966; Chapter IV, Section 20.0, pp 343−349. (43) Leah, R. T.; Brandon, N. P.; Vesovic, V.; Kelsall, G. H. Numerical Modeling of the Mass Transport and Chemistry of a Simplified Membrane-Divided Chlor-Alkali Reactor. J. Electrochem. Soc. 2000, 147, 4173. (44) Cheng, C. Y.; Kelsall, G. H. Models of Hypochlorite Production in Electrochemical Reactors with Plate and Porous Anodes. J. Appl. Electrochem. 2007, 37, 1203−1217. (45) Ping Ong, S.; Wang, L.; Kang, B.; Ceder, G. Li-Fe-P-O2 Phase Diagram from First Principles Calculations. Chem. Mater. 2008, 20, 1798−1807. (46) Jain, A.; Hautier, G.; Ong, S. P.; Moore, C. J.; Fischer, C. C.; Persson, K. A.; Ceder, G. Formation Enthalpies by Mixing GGA and GGA + U Calculations. Phys. Rev. B: Condens. Matter Mater. Phys. 2011, 84, 045115−1−045115−10. (47) Jain, A.; Ong, S. P.; Hautier, G.; Chen, W.; Richards, W. D.; Dacek, S.; Cholia, S.; Gunter, D.; Skinner, D.; Ceder, G. e.; Persson, K. A. Commentary: the Materials Project: a Materials Genome Approach to Accelerating Materials Innovation. APL Mater. 2013, 1, 011002−1− 011002−11. (48) Wagman, D. D.; Evans, W. H.; Parker, V. B.; Schumm, R. H.; Halow, I.; Bailey, S. M.; Churney, K. L.; Nuttall, R. L. The NBS Tables of Chemical Thermodynamic Properties: Selected Values for Inorganic and C1 and C2 Organic Substances in SI Units. J. Phys. Chem. Ref. Data, Suppl. 1982, 11 (2), 1−407. (49) Wendt, H., Kreysa, G. Electrochemical Engineering; Springer Science + Business Media, 1999. (50) van der Stegen, J. H. G. The State of the Art of Modern Chlor Alkali Electrolysis with Membrane Cells. Ph.D. Thesis, University of Twente, Holland, 2000. (51) Colman, J. E. Electrolytic Production of Sodium Chlorate. AIChE Symp. Ser. 1981, 77, 244−263. (52) De Valera, V. On the Theory of Electrochemical Chlorate Formation. Trans. Faraday Soc. 1953, 49, 1338−1351. (53) Landolt, D.; Ibl, N. On the Mechanism of Anodic Chlorate Formation in Concentrated NaCl Solutions. Electrochim. Acta 1970, 15, 1165−1183. (54) Hammar, L.; Wranglén, G. Cathodic and Anodic Efficiency Losses in Chlorate Electrolysis. Electrochim. Acta 1964, 9, 1−16. (55) Jaksić, M. M.; Despić, A.; Csonka, I.; Nikolić, B. Studies on Chlorate Cell Process V. Theory and Practice of a Modified Technology for Electrolytic Chlorate Production. J. Electrochem. Soc. 1969, 116, 1316−1322. (56) Jaksić, M. M. Mutual Effect of Current Density, pH, Temperature, and Hydrodynamic Factors on Current Efficiency in the Chlorate Cell Process. J. Electrochem. Soc. 1974, 121, 70−79. (57) Wanngård, J. Oxygen Formation in Chlorate Electrolysis. Poster presented at the Joint Meeting of Dechema and the Society of Chemical Industry on Electrochemical Cell Design and Optimization Procedures, 1990. (58) Foerster, F.; Müller, E. Beiträge zur Theorie der Elektrolyse von Alkalichloridlösungen. Z. Electrochem. 1903, 9, 171−185. 3023


Chemical Reviews

Review

(59) Foerster, F. The Electrolysis of Hypochlorite Solutions. Trans. Am. Electrochem. Soc. 1924, 46, 23−50. (60) Spasojević, M.; Ribić-Zelenović, L.; Spasojević, P. Microstructure of New Composite Electrocatalyst and Its Anodic Behavior for Chlorine and Oxygen Evolution. Ceram. Int. 2012, 38, 5827−5833. (61) Ibl, N.; Landolt, D. On The Mechanism of Anodic Chlorate Formation in Dilute NaCl solutions. J. Electrochem. Soc. 1968, 115, 713− 720. (62) Ibl, N.; Vogt, H. In Comprehensive Treatise of Electrochemistry; Bockris, J. O., Ed.; Plenum Press, 1981; Vol. 2, Chapter 3, pp 167−250. (63) Exner, K. S.; Anton, J.; Jacob, T.; Over, H. Controlling Selectivity in the Chlorine Evolution Reaction over RuO2-Based Catalysts. Angew. Chem., Int. Ed. 2014, 53, 11032−11035. (64) Karlsson, R. K.; Hansen, H. A.; Bligaard, T.; Cornell, A.; Pettersson, L. G. Ti Atoms in Ru0.3Ti0.7O2 Mixed Oxides Form Active and Selective Sites for Electrochemical Chlorine Evolution. Electrochim. Acta 2014, 146, 733−740. (65) Karlsson, R. K. B.; Cornell, A.; Pettersson, L. G. M. The Electrocatalytic Properties of Doped TiO2. Electrochim. Acta 2015, 180, 514−527. (66) Buné, N.; Shilyaeva, G.; Losev, V. Electrode Kinetics of Side Reactions Oxygen Evolution from Chloride Solutions at Titaniumruthenium Oxide Anodes. Elektrokhimiya 1977, 13, 1540−1546. (67) Pecherskii, M.; Gorodetskii, V.; Buné, N.; Losev, V. Kinetics of the Side Electrode Processes on Ruthenium-titanium Anodes in Chloride Solutions. Elektrokhimiya 1982, 18, 415−422. (68) Buné, N.; Losev, V. V.; Reznik, M. F.; Zaripova, E. Selectivity and Electrochemical Behavior of Titanium-ruthenium Oxide Anodes Containing Different Levels of Ruthenium Dioxide. Elektrokhimiya 1986, 22, 396−398. (69) Tilak, B. V.; Tari, K.; Hoover, C. Metal Anodes and Hydrogen Cathodes: Their Activity Towards O2 Evolution and ClO3−Reduction Reactions. J. Electrochem. Soc. 1988, 135, 1386−1392. (70) Kuhn, A. T.; Mortimer, C. J. The Efficiency of Chlorine Evolution in Dilute Brines on Ruthenium Dioxide Electrodes. J. Appl. Electrochem. 1972, 2, 283−287. (71) Spasojević, M.; Krstajić, N.; Jaksić, M. A Selective Catalyst for Titanium Anodes: Development and Optimization: II. Selectivity Features. J. Res. Inst. Catal., Hokkaido Univ. 1984, 32, 29−36. (72) Spasojević, M.; Krstajić, N.; Jaksić, M. Electrocatalytic Optimization of Faradaic Yields in the Chlorate Cell Process. Surf. Technol. 1984, 21, 19−26. (73) Krstajic, N.; Spasojević, M.; Jaksic, M. A Selective Catalyst for Titanium Anodes: Development and Optimization I. Catalyst Structure, Activity and Durability. J. Res. Inst. Catal., Hokkaido Univ. 1984, 32, 19− 28. (74) Eberil, V. I.; Fedotova, N. S.; Novikov, E. A.; Mazanko, A. F. Studying the Link Between the Potential of a Metal-oxide Anode, the Current Efficiency for Chlorate, and the Current Losses for the Oxygen and Chlorine Evolution in a Wide Range of the Chlorate Electrolysis Conditions. Russ. J. Electrochem. 2000, 36, 1296−1302. (75) Arikawa, T.; Murakami, Y.; Takasu, Y. Simultaneous Determination of Chlorine and Oxygen Evolving at RuO2/Ti and RuO2−TiO2/ Ti Anodes by Differential Electrochemical Mass Spectroscopy. J. Appl. Electrochem. 1998, 28, 511−516. (76) Jirkovský, J.; Hoffmannová, H.; Klementová, M.; Krtil, P. Particle Size Dependence of the Electrocatalytic Activity of Nanocrystalline RuO2 Electrodes. J. Electrochem. Soc. 2006, 153, E111−E118. (77) Macounová, K.; Makarova, M.; Jirkovský, J.; Franc, J.; Krtil, P. Parallel Oxygen and Chlorine Evolution on Ru1−xNixO2−y Nanostructured Electrodes. Electrochim. Acta 2008, 53, 6126−6134. (78) Makarova, M. V.; Jirkovský, J.; Klementová, M.; Jirka, I.; Macounová, K.; Krtil, P. The Electrocatalytic Behavior of Ru0.8Co0.2O2−x - the Effect of Particle Shape and Surface Composition. Electrochim. Acta 2008, 53, 2656−2664. (79) Macounová, K.; Jirkovský, J.; Makarova, M. V.; Franc, J.; Krtil, P. Oxygen Evolution on Ru1−xNixO2−y Nanocrystalline Electrodes. J. Solid State Electrochem. 2009, 13, 959−965.

(80) Macounová, K.; Makarova, M.; Krtil, P. Oxygen Evolution on Nanocrystalline RuO2 and Ru0.9Ni0.1O2−δ Electrodes - DEMS Approach to Reaction Mechanism Determination. Electrochem. Commun. 2009, 11, 1865−1868. (81) Petrykin, V.; Bastl, Z.; Franc, J.; Macounova, K.; Makarova, M.; Mukerjee, S.; Ramaswamy, N.; Spirovova, I.; Krtil, P. Local Structure of Nanocrystalline Ru1−xNixO2−δ Dioxide and Its Implications for Electrocatalytic Behaviors - an XPS and XAS Study. J. Phys. Chem. C 2009, 113, 21657−21666. (82) Petrykin, V.; Macounová, K.; Okube, M.; Mukerjee, S.; Krtil, P. Local Structure of Co Doped RuO2 Nanocrystalline Electrocatalytic Materials for Chlorine and Oxygen Evolution. Catal. Today 2013, 202, 63−69. (83) Halck, N. B.; Petrykin, V.; Krtil, P.; Rossmeisl, J. Beyond the Volcano Limitations in Electrocatalysis−Oxygen Evolution Reaction. Phys. Chem. Chem. Phys. 2014, 16, 13682−13688. (84) Macounová, K. M.; Simic, N.; Ahlberg, E.; Krtil, P. Electrochemical Water-Splitting Based on Hypochlorite Oxidation. J. Am. Chem. Soc. 2015, 137, 7262−7265. (85) Chapin, R. M. The Effect of Hydrogen-Ion Concentration on the Decomposition of Hypohalites. J. Am. Chem. Soc. 1934, 56, 2211−2215. (86) Lister, M. Decomposition of Sodium Hypochlorite: the Uncatalyzed Reaction. Can. J. Chem. 1956, 34, 465−478. (87) Bhaduri, I. Ü ber die Umsetzung der Hypochlorite in Chlorate. Z. Anorg. Chem. 1897, 13, 385−406. (88) Lister, M. W. The Decomposition of Hypochlorous Acid. Can. J. Chem. 1952, 30, 879−889. (89) Lister, M. W.; Petterson, R. C. Oxygen Evolution from Sodium Hypochlorite Solutions. Can. J. Chem. 1962, 40, 729−733. (90) Church, J. A. Kinetics of the Uncatalyzed and Copper(II)catalyzed Decomposition of Sodium Hypochlorite. Ind. Eng. Chem. Res. 1994, 33, 239−245. (91) Adam, L. C.; Gordon, G. Hypochlorite Ion Decomposition: Effects of Temperature, Ionic Strength, and Chloride Ion. Inorg. Chem. 1999, 38, 1299−1304. (92) Sandin, S.; Karlsson, R. K. B.; Cornell, A. The Catalyzed and Uncatalyzed Decomposition of Hypochlorite. Ind. Eng. Chem. Res. 2015, 54, 3767−3774. (93) Howell, O. R. The Catalytic Decomposition of Sodium Hypochlorite by Cobalt Peroxide. Proc. R. Soc. London, Ser. A 1923, 104, 134−152. (94) White, A. D. Action of Solutions of Bleaching Powder and of Hypochlorous acid on Metals. J. Soc. Chem. Ind. 1903, 22, 132−134. (95) Hofmann, K. A.; Ritter, K. Beständigkeit und Oxydationspotential der Hypochlorite, Beiträge zur Katalyse und über ein HypochloritKohle-Element. Ber. Dtsch. Chem. Ges. 1914, 47, 2233−2244. (96) Chirnoaga, E. CCXXII. - the Catalytic Decomposition of Solutions of Sodium Hypochlorite by Finely Divided Metallic Oxides. J. Chem. Soc. 1926, 129, 1693−1703. (97) Lewis, J. R. The Catalytic Decomposition of Sodium Hypochlorite Solutions. I. Mechanism of the Reaction. J. Phys. Chem. 1927, 32, 243−254. (98) Lewis, J. R. The Catalytic Decomposition of Sodium Hypochlorite Solutions. III. Promoter Action of Hydrated Magnaesium Oxide in the Hydrated Copper Oxide Catalysis of Sodium Hypochlorite. J. Phys. Chem. 1930, 35, 915−919. (99) Connick, R. E.; Hurley, C. R. Chemistry of Ru(VI), -(VII) and -(VIII). Reactions, Oxidation Potentials and Spectra. J. Am. Chem. Soc. 1952, 74, 5012−5015. (100) Patel, M. C.; Mankad, B. N. A Note on the Effect of Salts on the Decomposition of Sodium Hypochlorite. J. Indian Chem. Soc., Ind. News Ed. 1954, 17, 236−237. (101) Ayres, G. H.; Booth, M. H. Catalytic Decomposition of Hypochlorite Solution by Iridium Compounds. I. The pH-Time Relationship. J. Am. Chem. Soc. 1955, 77, 825−827. (102) Ayres, G. H.; Booth, M. H. Catalytic Decomposition of Hypochlorite Solution by Iridium Compounds. II. Kinetic Studies. J. Am. Chem. Soc. 1955, 77, 828−833. 3024


Chemical Reviews

Review

(103) Lister, M. Decomposition of Sodium Hypochlorite: the Catalyzed Reaction. Can. J. Chem. 1956, 34, 479−488. (104) Gray, E. T.; Taylor, R. W.; Margerum, D. W. Kinetics and Mechanisms of the Coppercatalyzed Decomposition of Hypochlorite and Hypobromite. Properties of a Dimeric Copper(III) Hydroxide Intermediate. Inorg. Chem. 1977, 16, 3047−3055. (105) Hamano, A.; Mori, H. Activity of Metal Oxide Catalysts. 4. Activity and Surface Area of Transition Metal Oxide Catalysts. Sasebo Kogyo Koto Senmon Gakko Kenkyu Hokoku 1977, 14, 65−72. (106) Hamano, A.; Takakura, K.; Nakamori, I. The Effect of pH on the Catalytic Decomposition of Aqueous Sodium Hypochlorite by Cobalt Oxide. Kayaku Gakkaishi 1998, 59, 174−180. (107) Baldwin, R. T. History of the Chlorine Industry. J. Chem. Educ. 1927, 4, 313−319. (108) Kershaw, J. B. C. The Present Position and Future Prospects of the Electrolytic Alkali and Bleach Industry. Trans. Faraday Soc. 1907, 3, 38−47. (109) Müller, E. Alkalichloridelektrolyse. Z. Elektrochem. Angew. Phys. Chem. 1907, 13, 800−804. (110) Oettel, F. Studien Ü ber Die Elektrolytische Bildung Von Unterchlorig-sauren Und Chlorsauren Salzen. Z. Elektrochem. Angew. Phys. Chem. 1894, 1, 354−361. (111) Oettel, F. Studien Ü ber Die Elektrolytische Bildung Von Unterchlorig-sauren Und Chlorsauren Salzen. II. Z. Elektrochem. Angew. Phys. Chem. 1895, 1, 474−480. (112) Oettel, F. Zur Elektrolyse von Chlorcalcium-Lösungen. Z. Elektrochem. Angew. Phys. Chem. 1898, 5, 1−5. (113) Bischoff, H.; Foerster, F. Beitrag Zur Kenntnis Der Elektrolyse Von Chlorcalciumlösungen. Z. Elektrochem. 1898, 4, 464−470. (114) Foerster, F.; Müller, E.; Jorre, F. Zur Kenntnis der Vorgänge bei der Elektrolyse der Alkalichloridlösungen. Z. Elektrochem. Angew. Phys. Chem. 1899, 6, 11−26. (115) Foerster, F.; Jorre, F. Zur Kenntniss der Beziehungen der unterchlorigsauren Salze zu den chlorsauren Salzen. J. Prakt. Chem. 1899, 59, 53−101. (116) Foerster, F.; Sonneborn, H. Zur Kenntnis der anodischen Sauerstoffentwicklung bei der Elektrolyse von Alkalichloridlösungen. Z. Elektrochem. Angew. Phys. Chem. 1900, 6, 597−604. (117) Foerster, F.; Jorre, F. Zur Kenntnis der Erscheinungen bei der Elektrolyse von Alkalichloridlösungen mit Diaphragma. Z. Anorg. Chem. 1900, 23, 158−219. (118) Foerster, F. Zur Theorie der elektrolytischen Bildung von Hypochlorit und Chlorat. Z. Anorg. Chem. 1899, 22, 1−32. (119) Foerster, F.; Müller, E. Zur Kenntnis der Elektrolyse, zumal der Alkalichloride, an platinierten Elektroden. Z. Elektrochem. Angew. Phys. Chem. 1902, 8, 515−540. (120) Foerster, F.; Müller, E. Zur Theorie der Einwirkung der Halogene auf Alkalien. Z. Elektrochem. Angew. Phys. Chem. 1902, 8, 921− 926. (121) Foerster, F.; Müller, E. Ü ber die bei der elektrolytischen Darstellung von Alkali- Hypochloriten und -Chloraten erreichbaren Strom- und Energie-Ausbeuten. Z. Elektrochem. Angew. Phys. Chem. 1902, 8, 8−17. (122) Foerster, F.; Müller, E. Ü ber das Verhalten der unterchlorigen Säure und ihrer Salze bei der Elektrolyse. Z. Elektrochem. Angew. Phys. Chem. 1902, 8, 633−638. (123) Foerster, F. Ü ber kÜ nstlichen Graphit und über Platiniridium als Anodenmaterialien. Z. Elektrochem. Angew. Phys. Chem. 1902, 8, 143− 147. (124) Foerster, F.; Dolch, P. Die Umwandlung von Hypochlorit in Chlorat in alkalischer Lösung. Z. Elektrochem. 1917, 23, 137−147. (125) Foerster, F.; Müller, E. Ü ber das Verhalten der unterchlorigen Säure und ihrer Salze bei der Elektrolyse. Z. Elektrochem. Angew. Phys. Chem. 1902, 8, 665−672. (126) Sproesser, L. Ü ber Alkalichlorid-Elektrolyse an Kohlenanoden. Z. Elektrochem. Angew. Phys. Chem. 1901, 7, 971−976. (127) Oechsli, W. Ü ber die Elektrolytische Perchloratbildung. Z. Elektrochem. Angew. Phys. Chem. 1903, 9, 807−828.

(128) Müller, E. Experimentaluntersuchung über die Bildung von Hypochlorit und Chlorat bei der Elektrolyse von Alkalichloriden. Z. Anorg. Chem. 1899, 22, 33−90. (129) Müller, E.; Koppe, P. Einfluss der Stromkonzentration auf die Elektrolytische Chloratbildung. Z. Elektrochem. 1911, 17, 421−430. (130) Knibbs, N.; Palfreeman, H. The Theory of Electrochemical Chlorate and Perchlorate Formation. Trans. Faraday Soc. 1920, 16, 402−433. (131) Joffe, W. S. Ü ber den Anodenprozess bei der Elektrolyse wässriger Alkalichloridlösungen mit imprägnierten und nichtimprägnierten Graphitelektroden. II: Gemeinsame Abscheidung von Chlor und Sauerstoff an der Anode. Zeitschr. Electrochem. 1936, 42, 76−79. (132) Rius, A.; Llopis, J. Sobre La Oxidacion Anodica De Los Hipocloritos Alcalinos. Anales De Fisica y Quimica 1945, 41, 1030−1053. (133) de Valera, V. The Mechanism of Chloride Electrolysis. Trans. Faraday Soc. 1956, 52, 250−260. (134) Despić, A.; Jaksić, M.; Nikolić, B. The Effect of Kinetic and Hydrodynamic Factors on Current Efficiency in the Chlorate Cell Process. J. Appl. Electrochem. 1972, 2, 337−343. (135) Jaksić, M. M.; Nikolić, B.; Csonka, I.; Djordjević, A. Studies on Chlorate Cell Process IV. Effect of Neutral Salts on Conversion of Available Chlorine to Chlorate. J. Electrochem. Soc. 1969, 116, 684−687. (136) Kuhn, A. T.; Mortimer, C. J. The Kinetics of Chlorine Evolution and Reduction on Titanium-SupportedMetal Oxides Especially RuO2 and IrO2. J. Electrochem. Soc. 1973, 120, 231−236. (137) Jaksić, M. M.; Despić, A.; Nikolić, B.; Maksić, S. Effect of Some Anions on the Chlorate Cell Process. Croat. Chem. Acta 1972, 44, 61− 66. (138) Bondar, R.; Borisova, A. A.; Kalinovskii, E. A. Titaniumruthenium Dioxide Anodes in the Electrolysis of Chloride-sulphate Solutions. Elektrokhimiya 1974, 10, 44−48. (139) Veselovskaya, I. E.; Spasskaya, E. K.; Sokolov, V. A.; Takachenko, V. I.; Yakimenko, L. M. Titanium-ruthenium Dioxide Anodes in the Electrolysis of Chloridesulphate Solutions. Elektrokhimiya 1974, 10, 70−73. (140) Agapova, R. A.; Kokhanov, G. N. Electrochemical Properties of Cobalt Oxide Anodes. Elektrokhimiya 1976, 12, 1649−1653. (141) Shlyapnikov, V. A. Role of Anode Material in Electrosynthesis of Chlorates. J. Appl. Chem. USSR 1976, 49, 90−94. (142) Uzbekov, A. A.; Lambrev, V. G.; Yazikov, I. F.; Buné, N. Y.; Gorodetskii, V. V.; Losev, V. V. Possibility of Improving Chloride Electrolysis by Use of Titanium-supported Ruthenium Dioxide Anodes. Elektrokhimiya 1978, 14, 1303. (143) Kazarinov, V. E.; Andreev, V. N. Investigation of the Adsorption of Ions and the Structure of the Electric Double Layer on Rutheniumtitanium Oxide Anodes. Elektrokhimiya 1978, 14, 577−579. (144) Fukuda, K.-i.; Iwakura, C.; Tamura, H. Effect of the Addition of NH4F on Anodic Behaviors of DSA-type Electrodes in H2SO4(NH4)2SO4 Solutions. Electrochim. Acta 1979, 24, 367−371. (145) Kelsall, G.; Robbins, D. Thermodynamics of Ti-H2O-F(-Fe) systems at 298 K. J. Electroanal. Chem. Interfacial Electrochem. 1990, 283, 135−157. (146) Iwakura, C.; Inai, M.; Tamura, H. Foreign Metal-doped SnO2 Film Anodes for Oxygen and Chlorine Evolution. Chem. Lett. 1979, 8, 225−228. (147) Caldwell, D. L.; Hazelrigg, M. J. In Modern Chlor-Alkali Technology; Coulter, M. O., Ed.; SCI, 1980; Vol. 1; Chapter 10, pp 137− 144. (148) Saito, S. In Modern Chlor-Alkali Technology; Coulter, M. O., Ed.; SCI, 1980; Vol. 1; Chapter 11, pp 137−144. (149) Harjanto, S.; Cao, Y.; Shibayama, A.; Naitoh, I.; Nanami, T.; Kasahara, K.; Okumura, Y.; Liu, K.; Fujita, T. Leaching of Pt, Pd and Rh from Automotive Catalyst Residue in Various Chloride Based Solutions. Mater. Trans. 2006, 47, 129−135. (150) Bergner, D. Stromausbeuten, Chlorat- und Sauerstoffbildung bei der Alkali-chlorid- Elektrolyse nach dem Membranverfahren. Chem.-Ztg. 1980, 104, 215−224. (151) Gorodetskii, V. V.; Pecherskii, M. M.; Yanke, V. B.; Buné, N. Y.; Busse-Machukas, V. B.; Kubasov, V. L.; Losev, V. V. Effect of Acidity on 3025


Chemical Reviews

Review

Mathemeatical and Computational Methods in Electrochemical Engineering; Electrochemical Society: Pennington, NJ, 1993; 147−154. (173) Baolian, Y.; Wenhua, Z. Current Efficiency in the Chlorate Cell Process with an Oxygen Cathode (III): a Gas Analysis Method for Chlorate Current. J. Appl. Electrochem. 1994, 24, 503−508. (174) Cornell, A.; Håkansson, B.; Lindbergh, G. Ruthenium Based DSA® in Chlorate Electrolysis - Critical Anode Potential and Reaction Kinetics. Electrochim. Acta 2003, 48, 473−481. (175) Gritzner, G.; Kreysa, G. Nomenclature, Symbols and Definitions in Electrochemical Engineering. Pure Appl. Chem. 1993, 65, 1009−1020. (176) Byrne, P.; Fontes, E.; Parhammar, O.; Lindbergh, G. A Simulation of the Tertiary Current Density Distribution from a Chlorate Cell: I. Mathematical Model. J. Electrochem. Soc. 2001, 148, D125−D132. (177) Byrne, P. Mathematical Modelling and Experimental Simulation of Chlorate and Chloralkali Cells. Ph.D. Thesis, KTH The Royal institute of technology, Stockholm, Sweden, 2001. (178) Santana, M. H.; De Faria, L. A. Oxygen and Chlorine Evolution on RuO2+TiO2+CeO2+ Nb2O5 Mixed Oxide Electrodes. Electrochim. Acta 2006, 51, 3578−3585. (179) Panić, V.; Jovanović, V.; Terzić, S.; Barsoum, M.; Jović, V.; Dekanski, A. The Properties of Electroactive Ruthenium Oxide Coatings Supported by Titanium-based Ternary Carbides. Surf. Coat. Technol. 2007, 202, 319−324. (180) Macounová, K.; Makarova, M.; Franc, J.; Jirkovský, J.; Krtil, P. Influence of Oxygen on Reactivity of Ru1−xFexO2−y-Doped Materials. Electrochem. Solid-State Lett. 2008, 11, F27−F29. (181) Petrykin, V.; Macounova, K.; Shlyakhtin, O.; Krtil, P. Tailoring the Selectivity for Electrocatalytic Oxygen Evolution on Ruthenium Oxides by Zinc Substitution. Angew. Chem., Int. Ed. 2010, 49, 4813− 4815. (182) Petrykin, V.; Macounova, K.; Franc, J.; Shlyakhtin, O.; Klementova, M. K.; Mukerjee, S.; Krtil, P. Zn-doped RuO2 Electrocatalysts for Selective Oxygen Evolution: Relationship Between Local Structure and Electrocatalytic Behavior in Chloride Containing Media. Chem. Mater. 2011, 23, 200−207. (183) Hansen, H. A.; Man, I. C.; Studt, F.; Abild-Pedersen, F.; Bligaard, T.; Rossmeisl, J. Electrochemical Chlorine Evolution at Rutile Oxide (110) Surfaces. Phys. Chem. Chem. Phys. 2010, 12, 283−290. (184) Rossmeisl, J.; Qu, Z.-W.; Zhu, H.; Kroes, G.-J.; Nørskov, J. K. Electrolysis of Water on Oxide Surfaces. J. Electroanal. Chem. 2007, 607, 83−89. (185) Chen, S.; Zheng, Y.; Wang, S.; Chen, X. Ti/RuO2-Sb2O5-SnO2 Electrodes for Chlorine Evolution from Seawater. Chem. Eng. J. 2011, 172, 47−51. (186) Wang, S.; Xu, H.; Yao, P.; Chen, X. Ti/RuO2-IrO2SnO2-Sb2O5 anodes for Cl2 Evolution from Seawater. Electrochemistry 2012, 80, 507− 511. (187) Neodo, S.; Rosestolato, D.; Ferro, S.; De Battisti, A. On the Electrolysis of Dilute Chloride Solutions: Influence of the Electrode Material on Faradaic Efficiency for Active Chlorine, Chlorate and Perchlorate. Electrochim. Acta 2012, 80, 282−291. (188) Xiong, K.; Deng, Z.; Li, L.; Chen, S.; Xia, M.; Zhang, L.; Qi, X.; Ding, W.; Tan, S.; Wei, Z. Sn and Sb Co-doped RuTi Oxides Supported on TiO2 Nanotubes Anode for Selectivity Toward Electrocatalytic Chlorine Evolution. J. Appl. Electrochem. 2013, 43, 847−854. (189) Kuznetsova, E.; Petrykin, V.; Sunde, S.; Krtil, P. Selectivity of Nanocrystalline IrO2- based Catalysts in Parallel Chlorine and Oxygen Evolution. Electrocatalysis 2015, 6, 198−210. (190) Le Luu, T.; Kim, J.; Yoon, J. Physicochemical Properties of RuO2 and IrO2 Electrodes Affecting Chlorine Evolutions. J. Ind. Eng. Chem. 2015, 21, 400−404. (191) Spasojevic, M.; Ribic-Zelenovic, L.; Spasojevic, P.; Nikolic, B. Current Efficiency In The Chlorate Cell Process. J. Serb. Chem. Soc. 2014, 79, 677−688. (192) Zeradjanin, A. R.; Menzel, N.; Schuhmann, W.; Strasser, P. On the Faradaic Selectivity and the Role of Surface Inhomogeneity During the Chlorine Evolution Reaction at Ternary Ti-Ru-Ir Mixed Metal

the Electrochemical and Corrosion Behavior of Titanium-ruthenium Oxide Anodes in Chloride Solutions. Elektrokhimiya 1981, 17, 513−517. (152) Érenburg, R. G.; Krishtalik, L. I.; Yaroshevskaya, I. P. Kinetics of chlorine evolution and ionization at ruthenium oxide and rutheniumtitanium oxide anodes. Elektrokhimiya 1975, 11, 1072−1074. (153) O’Grady, W.; Iwakura, C.; Huang, J.; Yeager, E. Report Number 37: Ruthenium Oxide Catalysts for the Oxygen Electrode; Office of Naval Research: Department of Chemistry, Case Western Reserve University: Cleveland, Ohio, 1974. (154) Denton, D. A.; De Souza, J. T.; Entwisle, J. H.; Lee, D.; Wilson, H. G. In Developments in Coatings for Metal Anodes; Jackson, C., Ed.; Ellis Horwood: Chichester, 1983; Vol. 2, Chapter 13. (155) Kokoulina, D.; Bunakova, L. Oxygen Evolution at the Oxide Electrodes RuO2 and RuO2+TiO2 in Chloride Solutions. J. Electroanal. Chem. Interfacial Electrochem. 1984, 164, 377−383. (156) Trasatti, S. Electrocatalysis in the Anodic Evolution of Oxygen and Chlorine. Electrochim. Acta 1984, 29, 1503−1512. (157) Buné, N. Y.; Portnova, M. Y.; Filatov, V. P.; Losev, V. V. Effect of Foreign Anions on the Kinetics of Chlorine and Oxygen Evolution on the Ruthenium-titanium Oxide Anodes Under the Conditions of Chlorine Evolution. Elektrokhimiya 1984, 20, 1291−1295. (158) Cairns, J. F.; Couper, A. M.; Denton, D. A. In Modern Chloralkali Technology; Wall, K., Ed.; Ellis Horwood Publishers, 1986; Vol. 3, Chapter 21, pp 293−308. (159) Kotowski, S.; Busse, B. In Modern Chlor-alkali Technology; Wall, K., Ed.; Ellis Horwood Publishers, 1986; Vol. 3, Chapter 22, pp 310− 323. (160) Wranglén, G. Kloratelektrolysens Teori. Tek. Tidskr. 1962, 10, 1−4. (161) Spasojević, M.; Krstajić, N.; Jakšić, M. Structure, Properties and Optimization of an Anodic Electrocatalyst: RuO2/TiO2 on Titanium. J. Mol. Catal. 1987, 40, 311−326. (162) Hardee, K. L.; Mitchell, L. K. The Influence of Electrolyte Parameters on the Percent Oxygen Evolved from a Chlorate Cell. J. Electrochem. Soc. 1989, 136, 3314−3318. (163) Zhinkin, N. V.; Novikov, E. A.; Fedotova, N. S.; Éberil, V. I.; Busse-Macukas, V. B. Oxygen Evolution and Ruthenium Losses from Active Coatings of Titanium-ruthenium Oxide Anodes in Sodium Chloride Solutions at Different Current Densities, Ph and Temperatures. Elektrokhimiya 1989, 25, 1094−1099. (164) Eberil, V. I.; Dobrov, Y. G.; Novikov, E. A.; Fedotova, N. S. The Current Efficiency for Chlorate and Current Losses Due to Oxygen Formation as a Function of Anodic Potential in the Electrolysis of Chloride-chlorate Solutions at RTO Anodes. Russ. J. Electrochem. 1997, 33, 570−572. (165) Trasatti, S.; Lodi, G. In Electrodes of Conductive Metallic Oxides. Part B; Trasatti, S., Ed.; Elsevier Scientific Publishing Company, 1981; Chapter 10, pp 521−614. (166) Boodts, J. F. C.; Trasatti, S. Effect of Composition on the Electrocatalytic Activity of the Ternary Oxide Ru0:3Ti(0:7−x)SnxO2. J. Electrochem. Soc. 1990, 137, 3784−3789. (167) Couper, A.; Brooks, W.; Denton, D. In Modern Chlor-alkali Technology; Prout, N., Moorhouse, J., Eds.; Elsevier Applied Science, 1990; Vol. 4, Chapter 8, pp 71−83. (168) Wanngård, J. In Modern Chlor-alkali Technology; Wellington, T. C., Ed.; Elsevier Applied Science, 1992; Vol. 5; Chapter 25, pp 295−306. (169) Traini, C.; Meneghini, G. In Modern Chlor-alkali technology; Wellington, T. C., Ed.; Elsevier Applied Science, 1992; Vol. 5; Chapter 23, pp 269−280. (170) Czarnetzki, L.; Janssen, L. Formation of Hypochlorite, Chlorate and Oxygen During NaCl Electrolysis from Alkaline Solutions at an RuO2/TiO2 Anode. J. Appl. Electrochem. 1992, 22, 315−324. (171) Takasu, Y.; Arikawa, T.; Sunohara, S.; Yahikozawa, K. Direct Detection of Competitively Evolving Chlorine and Oxygen at Anodes. J. Electroanal. Chem. 1993, 361, 279−281. (172) Nakajima, Y.; Nakamatsu, S.; Shimamune, T.; Matsuda, Y. Oxide-Coated Titanium Anode for Chlor-alkali Industry. Proceedings of the Symposia on Chlor-Alkali and Chlorate Production and New 3026


Chemical Reviews

Review

Oxide Electrocatalysts. Phys. Chem. Chem. Phys. 2014, 16, 13741− 13747. (193) Abbott, D. F.; Petrykin, V.; Okube, M.; Bastl, Z.; Mukerjee, S.; Krtil, P. Selective Chlorine Evolution Catalysts Based on Mg-doped Nanoparticulate Ruthenium Dioxide. J. Electrochem. Soc. 2015, 162, H23−H31. (194) Calle-Vallejo, F.; Koper, M. T. First-principles Computational Electrochemistry: Achievements and Challenges. Electrochim. Acta 2012, 84, 3−11. (195) Nørskov, J. K.; Rossmeisl, J.; Logadottir, A.; Lindqvist, L.; Kitchin, J. R.; Bligaard, T.; Jónsson, H. Origin of the Overpotential for Oxygen Reduction at a Fuel-cell Cathode. J. Phys. Chem. B 2004, 108, 17886−17892. (196) Filhol, J.-S.; Neurock, M. Elucidation of the Electrochemical Activation of Water Over Pd by First Principles. Angew. Chem., Int. Ed. 2006, 45, 402−406. (197) Otani, M.; Sugino, O. First-Principles Calculations of Charged Surfaces and Interfaces: a Plane-wave Nonrepeated Slab Approach. Phys. Rev. B: Condens. Matter Mater. Phys. 2006, 73, 115407−1−115407−11. (198) Rossmeisl, J.; Skúlason, E.; Björketun, M. E.; Tripkovic, V.; Nørskov, J. K. Modeling the Electrified Solid-liquid Interface. Chem. Phys. Lett. 2008, 466, 68−71. (199) Nørskov, J. K.; Bligaard, T.; Logadottir, A.; Bahn, S.; Hansen, L.; Bollinger, M.; Bengaard, H.; Hammer, B.; Sljivancanin, Z.; Mavrikakis, M.; et al. Universality in Heterogeneous Catalysis. J. Catal. 2002, 209, 275−278. (200) Skúlason, E.; Karlberg, G. S.; Rossmeisl, J.; Bligaard, T.; Greeley, J.; Jónsson, H.; Nørskov, J. K. Density Functional Theory Calculations for the Hydrogen Evolution Reaction in an Electrochemical Double Layer on the Pt(111) Electrode. Phys. Chem. Chem. Phys. 2007, 9, 3241− 3250. (201) Skúlason, E.; Tripkovic, V.; Björketun, M. E.; Gudmundsdóttir, S.; Karlberg, G.; Rossmeisl, J.; Bligaard, T.; Jónsson, H.; Nørskov, J. K. Modeling the Electrochemical Hydrogen Oxidation and Evolution Reactions on the Basis of Density Functional Theory Calculations. J. Phys. Chem. C 2010, 114, 18182−18197. (202) Greeley, J.; Stephens, I. E. L.; Bondarenko, A. S.; Johansson, T. P.; Hansen, H. A.; Jaramillo, T. F.; Rossmeisl, J.; Chorkendorff, I.; Nørskov, J. K. Alloys of Platinum and Early Transition Metals as Oxygen Reduction Electrocatalysts. Nat. Chem. 2009, 1, 552−556. (203) Siahrostami, S.; Verdaguer-Casadevall, A.; Karamad, M.; Deiana, D.; Malacrida, P.; Wickman, B.; Escudero-Escribano, M.; Paoli, E. A.; Frydendal, R.; Hansen, T. W.; et al. Enabling Direct H2O2 Production Through Rational Electrocatalyst Design. Nat. Mater. 2013, 12, 1137− 1143. (204) Over, H.; Kim, Y. D.; Seitsonen, A. P.; Wendt, S.; Lundgren, E.; Schmid, M.; Varga, P.; Morgante, A.; Ertl, G. Atomic-Scale structure and Catalytic reactivity of the RuO2 (110) surface. Science 2000, 287, 1474− 1476. (205) Mars, P.; van Krevelen, D. Oxidations Carried Out by Means of Vanadium Oxide Catalysts. Chem. Eng. Sci. 1954, 3, 41−59. (206) Fang, Y.-H.; Liu, Z.-P. Mechanism and Tafel Lines of Electrooxidation of Water to Oxygen on RuO2(110). J. Am. Chem. Soc. 2010, 132, 18214−18222. (207) Exner, K. S.; Anton, J.; Jacob, T.; Over, H. Chlorine Evolution Reaction on RuO2(110): Ab initio Atomistic Thermodynamics Study Pourbaix Diagrams. Electrochim. Acta 2014, 120, 460−466. (208) Wang, H.; Schneider, W. F.; Schmidt, D. Intermediates and Spectators in O2 Dissociation at the RuO2 (110) Surface. J. Phys. Chem. C 2009, 113, 15266−15273. (209) Siahrostami, S.; Vojvodic, A. Influence of Adsorbed Water on the Oxygen Evolution Reaction on Oxides. J. Phys. Chem. C 2015, 119, 1032−1037. (210) Exner, K. S.; Anton, J.; Jacob, T.; Over, H. Microscopic Insights into the Chlorine Evolution Reaction on RuO2(110): a Mechanistic Ab Initio Atomistic Thermodynamics Study. Electrocatalysis 2015, 6, 163− 172. (211) Exner, K. S.; Anton, J.; Jacob, T.; Over, H. Ligand Effects and Their Impact on Electrocatalytic Processes Exemplified with the Oxygen

Evolution Reaction (OER) on RuO 2 (110). ChemElectroChem 2015, 2, 707−713. (212) Kötz, R.; Stucki, S.; Scherson, D.; Kolb, D. In-situ Identification of RuO4 as the Corrosion Product During Oxygen Evolution on Ruthenium in Acid Media. J. Electroanal. Chem. Interfacial Electrochem. 1984, 172, 211−219. (213) De Zoubov, N.; Pourbaix, M. In Atlas of Electrochemical Equilibria In Aqueous Solutions; Pourbaix, M., Ed.; Pergamon Press Ltd., 1966; Chapter IV, Section 13.1, pp 343−349. (214) Loučka, T. The Potential-pH Diagram for the Ru-H2O-Cl− System at 25°c. J. Appl. Electrochem. 1990, 20, 522−523. (215) Llopis, J.; Vázquez, M. Passivation of Ruthenium in Hydrochloric Acid Solution. Electrochim. Acta 1966, 11, 633−640. (216) Llopis, J.; Gamboa, J.; Alfayate, J. Radiochemical Study of the Anodic Corrosion of Ruthenium. Electrochim. Acta 1967, 12, 57−65. (217) Gorodetskii, V. V.; Pecherskii, M. M.; Skuratnik, Y. B.; Dembrovskii, M. A.; Losev, V. V. Anodic Behavior of Ruthenium. Elektrokhimiya 1973, 9, 894−897. (218) Pecherskii, M. M.; Gorodetskii, V. V.; Pulina, V. M.; Losev, V. V. Effect of Acidity on the Anodic Behavior of Ruthenium. Elektrokhimiya 1976, 12, 1445−1448. (219) Veselovskaya, I. E.; Khodkevich, S. D.; Malkina, R. I.; Yakimenko, L. M. Electrochemical Behavior of Ruthenium Anode in Chlorine and Oxygen Evolution Processes. Elektrokhimiya 1974, 10, 74−77. (220) Iwakura, C.; Hirao, K.; Tamura, H. Anodic Evolution of Oxygen on Ruthenium in Acidic Solutions. Electrochim. Acta 1977, 22, 329−334. (221) Uzbekov, A. A.; Lambrev, V. G.; Yazikov, I. F.; Rodin, N. N.; Zabrodskaya, L. M.; Klement’eva, V. S.; Vlodov, Y. M. Corrosion of Titanium-supported Ruthenium Dioxide Anodes - Nature of Anode Consumption and Ruthenium Dissolution Kinetics in Chloride Solutions. Elektrokhimiya 1978, 14, 1150−1159. (222) Kokoulina, D. V.; Krasovitskaya, Y. I.; Ivanova, T. V. Oxygen Evolution at Ruthenium Oxide Anodes and Its Connection with Anode Destruction. Elektrokhimiya 1978, 14, 470−474. (223) Bondar, R. U.; Kalinovskii, E. A. Electrochemical Stability of Titanium-ruthenium Oxide Anodes. Elektrokhimiya 1978, 14, 730−733. (224) Gorodetskii, V. V.; Pecherskii, M. M.; Yanke, V. B.; Shub, D. M.; Losev, V. V. Kinetics of Dissolution of Ruthenium-titanium Oxide Anodes in the Electrolysis of Chloride Solutions. Elektrokhimiya 1979, 15, 559−562. (225) Barral, G.; Guitton, J.; Montella, C.; Vergara, F. Étude des anodes à couche d’oxyde de ruthénium sur substrat de titane III: Influence de quelques paramétres de préparation des couches d’oxyde et des conditions d’électrolyse sur l’usure lors d’un dégagement d’oxygéne. Surf. Technol. 1980, 10, 25−45. (226) Tilak, B. V.; Birss, V. I.; Wang, J.; Chen, C.-P.; Rangarajan, S. K. Deactivation of Thermally Formed Ru/Ti Oxide Electrodes: An AC Impedance Characterization Study. J. Electrochem. Soc. 2001, 148, D112. (227) Eberil’, V. I.; Novikov, E. A.; Mazanko, A. F. Reasons for the DSA Passivation during Chlorate Electrolysis and the Means for Extending the Anode Service Life. Russ. J. Electrochem. 2001, 37, 1054−1058. (228) Burke, L. D.; Lyons, M. E.; McCarthy, M. Oxygen Evolution and Corrosion at Ruthenium Dioxide-Based Anodes. Proceedings 4th World Energy Conference, Pasadena, CA, 1982; IEEE: Piscataway, NJ, pp 267− 278. (229) Uzbekov, A. A.; Klement’eva, V. S. Radiochemical Investigation of the Selectivity of Dissolution of Components of Ruthenium-titanium Oxide Anodes (ORTA) in Chloride Solutions. Elektrokhimiya 1985, 21, 758−763. (230) Schmets, J.; Van Muylder, J.; Pourbaix, M. In Atlas of Electrochemical Equilibria In Aqueous Solutions; Pourbaix, M., Ed.; Pergamon Press Ltd, 1966; Chapter IV, pp 213−222. (231) Uzbekov, A. A.; Klement’eva, V. S.; Kubasov, V. L.; Lambrev, V. G. Corrosion Behavior of Ruthenium Oxide-titanium Anodes During Production of Sodium Chlorate by Electrolysis of Chloride-chlorate Solutions. J. Appl. Chem. USSR 1985, 58, 686−688. 3027


Chemical Reviews

Review

(232) Kötz, R.; Stucki, S. Stabilization of RuO2 by IrO2 for Anodic Oxygen Evolution in Acid Media. Electrochim. Acta 1986, 31, 1311− 1316. (233) Kolotyrkin, Y. M. Relationship Between Kinetic and Corrosion Characteristics of Anodic Materials. Indian J. Technol. 1986, 24, 381− 387. (234) Lyons, M. E. G.; Burke, L. D. Mechanism of Oxygen Reactions at Porous Oxide Electrodes. Part 1. - Oxygen Evolution at RuO2 and RuxSn1−xO2 Electrodes in Alkaline Solution Under Vigorous Electrolysis Conditions. J. Chem. Soc., Faraday Trans. 1 1987, 83, 299−321. (235) Manli, S.; Yanxi, C. In Modern Chlor-alkali Technology; Prout, N., Moorhouse, J., Eds.; Springer: The Netherlands, 1990; Vol. 4; Chapter 14, pp 149−157. (236) Hardee, K. L.; Kus, R. A. In Modern Chlor-alkali Technology; Sealey, S., Ed.; SCI, 1998; Vol. 7, Chapter 7, pp 43−54. (237) Bommaraju, T. V.; Chen, C.-P.; Birss, V. I. In Modern Chlor-alkali Technology; Moorhouse, J., Ed.; Wiley-Blackwell, 2001; Vol. 8, Chapter 5, pp 57−81. (238) Nylén, L.; Cornell, A. Critical Anode Potential in the Chlorate Process. J. Electrochem. Soc. 2006, 153, D14−D20. (239) Gajić-Krstajić, L.; Trišović, T.; Krstajić, N. Spectrophotometric Study of the Anodic Corrosion of Ti/RuO2 Electrode in Acid Sulfate Solution. Corros. Sci. 2004, 46, 65−74. (240) Cherevko, S.; Zeradjanin, A. R.; Topalov, A. A.; Kulyk, N.; Katsounaros, I.; Mayrhofer, K. J. J. Dissolution of Noble Metals during Oxygen Evolution in Acidic Media. ChemCatChem 2014, 6, 2219−2223. (241) Cherevko, S.; Reier, T.; Zeradjanin, A. R.; Pawolek, Z.; Strasser, P.; Mayrhofer, K. J. Stability of Nanostructured Iridium Oxide Electrocatalysts During Oxygen Evolution Reaction in Acidic Environment. Electrochem. Commun. 2014, 48, 81−85. (242) Zeradjanin, A. R.; Topalov, A. A.; Van Overmeere, Q.; Cherevko, S.; Chen, X.; Ventosa, E.; Schuhmann, W.; Mayrhofer, K. J. J. Rational Design of the Electrode Morphology for Oxygen Evolution - Enhancing the Performance for Catalytic Water Oxidation. RSC Adv. 2014, 4, 9579−9587. (243) Reier, T.; Pawolek, Z.; Cherevko, S.; Bruns, M.; Jones, T.; Teschner, D.; Selve, S.; Bergmann, A.; Nong, H. N.; Schlögl, R. e.; et al. Molecular Insight in Structure and Activity of Highly Efficient, Low-Ir Ir-Ni Oxide Catalysts for Electrochemical Water Splitting (OER). J. Am. Chem. Soc. 2015, 137, 13031−13040. (244) Topalov, A. A.; Katsounaros, I.; Auinger, M.; Cherevko, S.; Meier, J. C.; Klemm, S. O.; Mayrhofer, K. J. J. Dissolution of Platinum: Limits for the Deployment of Electrochemical Energy Conversion? Angew. Chem., Int. Ed. 2012, 51, 12613−12615. (245) Frydendal, R.; Paoli, E. A.; Knudsen, B. P.; Wickman, B.; Malacrida, P.; Stephens, I. E. L.; Chorkendorff, I. Benchmarking the Stability of Oxygen Evolution Reaction Catalysts: The Importance of Monitoring Mass Losses. ChemElectroChem 2014, 1, 2075−2081. (246) Frydendal, R.; Paoli, E. A.; Chorkendorff, I.; Rossmeisl, J.; Stephens, I. E. L. Toward an Active and Stable Catalyst for Oxygen Evolution in Acidic Media: Ti-Stabilized MnO2. Adv. Energy Mater. 2015, 5, 1−9. (247) Gijsbers, H. F. M.; Janssen, L. J. J. Distribution of mass transfer over a 0.5-m-tall hydrogen-evolving electrode. J. Appl. Electrochem. 1989, 19, 637−648. (248) Vogt, H.; Stephan, K. Local microprocesses at gas-evolving electrodes and their influence on mass transfer. Electrochim. Acta 2015, 155, 348−356. (249) Comninellis, C.; Vercesi, G. P. Problems in DSA® Coating Deposition by Thermal Decomposition. J. Appl. Electrochem. 1991, 21, 136−142. (250) Jones, G.; Bligaard, T.; Abild-Pedersen, F.; Nørskov, J. K. Using Scaling Relations to Understand Trends in the Catalytic Activity of Transition Metals. J. Phys.: Condens. Matter 2008, 30, 064239−1− 064239−6. (251) Rossmeisl, J.; Logadottir, A.; Nørskov, J. K. Electrolysis of Water on (Oxidized) Metal Surfaces. Chem. Phys. 2005, 319, 178−184. (252) Gorodetskii, V. V.; Zorin, P. N.; Pecherskii, M. M.; BusseMachukas, V. B.; Kubasov, V. L.; Tomashpol’skii, Y. Y. Auger

Spectroscopic Surface Composition Studies of Titanium-ruthenium Oxide Anodes (ORTA) Under the Conditions of Chlorine Electrolysis. Elektrokhimiya 1981, 17, 79−83. (253) Takasu, Y.; Sugimoto, W.; Nishiki, Y.; Nakamatsu, S. Structural analyses of RuO2−TiO2/Ti and IrO2−RuO2−TiO2/Ti anodes used in industrial chlor-alkali membrane processes. J. Appl. Electrochem. 2010, 40, 1789−1795. (254) Loučka, T. The Reason for the Loss of Activity of Titanium Anodes Coated with a Layer of RuO2 and TiO2. J. Appl. Electrochem. 1977, 7, 211−214.

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