Article pubs.acs.org/IECR
Catalyzed and Uncatalyzed Decomposition of Hypochlorite in Dilute Solutions Staffan Sandin, 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 S Supporting Information *
ABSTRACT: Hypochlorite decomposition has been investigated by the combined measurement of aqueous concentrations of total hypochlorite, chlorate, and chloride, as well as that of evolved oxygen. In all experiments, the initial concentrations of NaOCl and NaCl were 80 mM, and the temperature was 80 °C. The pH was kept constant in the range 5−10.5. The uncatalyzed decomposition of hypochlorite and the formation of chlorate and oxygen were all found to be third order of the form ri = ki[HOCl]2[OCl−], and kO2 was determined to be 0.046 M−2 s−1. A reaction mechanism in which oxygen and chlorate formation share an intermediate is proposed. Several compounds were tested for catalytic effects. The addition of chloride salts of cobalt and iridium showed catalytic effects on oxygen formation. The addition of iridium chloride also catalyzed the formation of chlorate with increasing selectivity for chlorate with increasing pH.
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INTRODUCTION Sodium chlorate (NaClO3) is produced industrially by electrolysis in undivided cells. The production amounts to over 3 million metric tons/year.1 Most sodium chlorate is used for bleaching of chemical pulp. The current efficiency for chlorate production is ∼95% and is limited by formation of oxygen.2 Improving the current efficiency by reducing oxygen formation could yield large economic savings due to the high yearly production of sodium chlorate. The production of sodium chlorate proceeds as follows. First, chlorine is produced on an anode from which it diffuses and is hydrolyzed to form hypochlorite (HOCl + OCl−) according to reactions 1 and 2 below. Hypochlorite further reacts to form chlorate according to reaction 3, which proceeds at its maximum rate when the ratio [HOCl]:[OCl−] equals 2. This corresponds to a pH of the electrolyte of 5.8−6.53, and therefore, the pH of chlorate electrolyte is maintained at 6−6.5. Cl 2 + H 2O ⇌ HOCl + Cl− + H+
(1)
HOCl ⇌ OCl− + H+
(2)
2HOCl + OCl− → ClO−3 + 2H+ + 2Cl−
(3)
costs for electricity has triggered the development of electrodes with increased activity and selectivity for both anode and cathode reactions. However, corrosion of such activated electrodes could release metallic compounds that may catalyze hypochlorite decomposition to oxygen in the electrolyte. It is therefore important to test whether potential electrode components can catalyze the decomposition of HOCl/OCl−. Hypochlorite decomposition is also a concern in chlor-alkali electrolysis, particularly in membrane cells operating with no anolyte acidification. The cation-selective membranes run at 97−98% efficiency. The remaining part of the current is transported by the migration of hydroxide ions to the anolyte. This causes an increase of the pH in the anolyte, leading to the formation of hypochlorite and subsequently to unwanted byproducts of oxygen in the cell gas and chlorate in the anolyte. The increase in pH can be counteracted by careful addition of HCl to the anolyte. Plants operating without anolyte acidification run at an anolyte pH > 4 with 1.5−2% oxygen in the chlorine gas and produce 3.5−5.2 kg ClO3−/ton NaOH.4,5 The goal of the present study is to clarify the processes of both uncatalyzed and catalyzed hypochlorite decomposition under conditions (pH, temperature, and hypochlorite concentration) relevant to the industrial chlorate process. However, industrial electrolytes also contain sodium chloride and sodium chlorate in high concentrations, typically 50−110 g/L of NaCl and 500−700 g/L of NaClO3.2 Nevertheless, we believe that the results, even though they should be applied with care to an industrial scale system, are useful for both chlorate producers, who could use the knowledge to guide which contaminants must be prevented from accumulating in the process electro-
Oxygen is the major byproduct in chlorate electrolysis and may be formed in anodic side reactions as well as in heterogeneous and homogeneous decomposition of hypochlorite, commonly presented as reactions 4 and 5 below. 2HOCl → O2 + 2H+ + 2Cl−
(4)
2OCl− → O2 + 2Cl−
(5)
Hydrogen gas is produced on the cathode side, and oxygen in the cell gas thus constitutes not only a loss in current efficiency but also a safety hazard, as explosive gas mixtures may form. Chlorate electrolysis requires large amounts of electrical energy, making up to 70% of the production cost. Increasing © 2015 American Chemical Society
Received: Revised: Accepted: Published: 3767
December 16, 2014 March 8, 2015 March 27, 2015 March 27, 2015 DOI: 10.1021/ie504890a Ind. Eng. Chem. Res. 2015, 54, 3767−3774
Article
Industrial & Engineering Chemistry Research
AgCl were weighed and added to the reactor directly. Reverse osmosis-filtered water (Milli-Q) was used in all trials and for diluting reagents. Instrumentation and Analytical Methods. The experimental setup consisted of a jacketed glass reaction vessel with temperature controlled by an external heater. The reaction mixture was continuously stirred by a magnetic stirrer. The pH was controlled by the addition of HCl and NaOH using an automatic titrator (Metrohm Titrando 907) with a combination glass electrode (Metrohm Unitrode with Pt 1000). The formed oxygen gas was measured with a mass spectrometer (Hiden Analytical HPR20). A known flow of argon, which was fed through the reactor, was used as a reference to determine the flow rate of oxygen at the detector. The mass spectrometer was calibrated prior to a run by measuring an air/argon mixture of known composition (0.9, 3.2, and 95.9% of O2, N2, and Ar, respectively). Samples were taken during the reaction at determined time intervals, typically, 1 and 7 min after hypochlorite addition and 1, 3, 5, 10, 15, 30, 60, and 90 min after acid addition. The samples were immediately stabilized by reduction of the hypochlorite with arsenite, resulting in the formation of arsenate and chloride.19 The concentrations of arsenate (which are directly proportional to the hypochlorite concentration in the reactor), chlorate, and chloride in the stabilized sample solutions were measured by ion chromatography (IC) using Transgenomic guard and separation columns and a membrane suppressor from Sequant. More detailed information regarding the IC equipment can be found in the Supporting Information. Between each experiment, the components that had been in contact with the reaction mixture were submerged in dilute nitric acid for at least 12 h and rinsed carefully prior to use. The experimental setup is shown in Figure 1.
lyte, and electrode manufacturers, who can use the information to guide which dopants might be unsuitable in chlorate electrodes. The uncatalyzed decomposition of hypochlorite at pH >9 has been shown to be of second order, as has chlorate and oxygen formation,6−8 whereas in the pH range of 5−8, the decomposition of hypochlorite has been found to be a third order reaction.9 Most studies on catalytic hypochlorite decomposition to oxygen have been made at alkaline conditions and at temperatures well below the operating temperature of modern chlorate cells.10−12 Furthermore, some studies made at pH 11.7 In the first work, a first order reaction mechanism for oxygen formation involving both HOCl and OCl−, with H2O2 as an intermediate, was proposed. In the second work, the oxygen formation was re-examined, and it was then described using a second order rate expression. This has been confirmed by Adam et al.6 Lister proposed an intermediate that decomposes into either oxygen or chlorite (ClO2−), which reacts further with hypochlorite to form chlorate. The intermediate proposed was Cl2O22−. This intermediate, although in a protonized form, has also been proposed by Adam et al.9 for the chlorate formation mechanism in the pH range 5−8. The initial steps of the mechanism proposed by Adam et al.9 can be seen below in reactions 12−14. 2HOCl ⇌ Cl 2O ·H 2O
HOCl + OCl− → O2 + 2Cl− + H+
Table 2. Compounds Tested, Their Catalytic Effects, the pH Values at which the Experiments Were Performeda
(12)
followed by OCl− + Cl 2O· H 2O → HOCl + HCl 2O2−
catalytic effects
(13)
or HOCl + Cl 2O·H 2O → HOCl + H 2Cl 2O2
(14)
In the case of chlorate formation, this intermediate decomposes to chlorite, which then reacts further to chlorate in a series of rapid reaction steps.9 We suggest that this intermediate may also, as proposed by Lister for Cl2O22−,7 decompose to oxygen instead of chlorite (eqs 15 and 16). Only a small fraction of the amount of intermediate formed decomposes to oxygen, and the majority reacts to form chlorate ions. HCl 2O2− → O2 + 2Cl− + H+ −
H 2Cl 2O2 → O2 + 2Cl + 2H
+
(15)
compound
O2
ClO3−
pH tested
AgCl Al2O3 CeCl3·7H2O CoCl2 Fe3O4 FeCl3·6H2O IrCl3·xH2O Na2Cr2O7·2H2O Na2MoO4·2H2O RuCl3·xH2O RuO2·xH2O
no no no yes no no yes no no no no
no no no no no no yes no no no no
6.5 6.5 6.5 6.5 6.5 6.5 6.5, 10.5 6.5, 11 6.5 6.5 6.5, 11.5
a
In all cases, except for Al2O3 (9 ppm) and AgCl (100 ppm), enough catalyst was added to achieve a final concentration of 10 μM in the reactor.
(16) 3771
DOI: 10.1021/ie504890a Ind. Eng. Chem. Res. 2015, 54, 3767−3774
Article
Industrial & Engineering Chemistry Research Catalytic effects on the rates were concluded if the rate of oxygen or chlorate formation, as estimated by the flow rate at the MS detector in the case of oxygen and by concentration changes in the case of chlorate, was significantly larger than the uncatalyzed rate. It is clear that the majority of the compounds do not catalyze the decomposition of dilute hypochlorite at 80 °C and neutral pH. An obvious difference between these experiments and the conditions under industrial chlorate production is that the concentrations of sodium chlorate and sodium chloride are much lower in these experiments. This could limit the oxidation states accessible for the metal compounds in these experiments. Nevertheless, it seems that none of the steel-alloying metals or Al2O3 have any effect on catalysis. Furthermore, Na2Cr2O7 does not catalyze either oxygen evolution or chlorate formation. AgCl does not have an effect either, despite the catalytic effect of silver in chlorine dioxide production,17 nor does the important electrocatalyst element Ru. However, both Co and Ir affect the reaction rates, and their effects will be discussed in more detail below. The catalytic effects on the formation of oxygen with Co and Ir, and the lack of a catalytic effect with Fe and Ru, agree with observations of Wanngård,24 who performed experiments using an industrial chlorate electrolyte. Cobalt Chloride. Cobalt is a known catalyst for the formation of oxygen from hypochlorite decomposition, and the influence of cobalt oxide has been studied by several groups.8,10−12 Furthermore, it has been studied as a possible dopant in electrocatalysts for chlorine evolution.15 In this study, cobalt was added as CoCl2. The addition had a catalytic effect on oxygen formation. Furthermore, a precipitate could be seen during the reaction. This precipitate is likely to be a cobalt oxide,18 and the oxygen formation reaction could therefore be heterogeneously catalyzed. However, the precipitate was not analyzed and conclusions regarding the catalytic mechanism can therefore not be made without further studies. Fitting of data from the cobalt-catalyzed decomposition of hypochlorite and formation of chlorate to a third order rate expression (7) yields the rate constants 3.8 M −2 s −1 for the total decomposition and 0.6 M−2 s−1 for chlorate formation. The third order rate constant for the oxygen formation was found to be 0.2 M−2 s−1. The fit was not as good as that to the uncatalyzed data. This can be seen by calculating the ratio of 3kClO3− + 2kO2 and kCH, which is 0.6. After the reaction, 0.23 mol of chlorate and 0.1 mol of oxygen had been formed per mole of reacted hypochlorite. Addition of cobalt salt was also seen to catalyze the formation of oxygen at pH levels of approximately 10 and 3, which agrees with the results of Hamano et al.8 Iridium Chloride. Although not as well-studied as cobalt oxides, iridium oxide has been shown to catalyze both chlorate and oxygen formation reactions.25,26 The same effects were observed in the current work, where iridium was added as IrCl3 to a concentration of 10 μM at 80 °C. This was done at both pH 6.5 and 10.5. In Figures 7 and 8, trials with iridium addition at pH 6.5 and 10.5 are compared to the uncatalyzed decomposition. Analysis of the results show that 0.18 mol of chlorate and 0.15 mol of oxygen were formed per reacted mole of hypochlorite at pH 6.5. At pH 10.5, 0.27 mol of chlorate and 0.04 mol of oxygen were formed per reacted mole of hypochlorite. This increased selectivity for chlorate formation with increasing pH was also noted by Ayres and Booth.25,26 As in the cobalt trial, a precipitate was observed, which was likely iridium oxide.18 The precipitate was not analyzed and definite
Figure 7. Measured concentrations in a reactor from IrCl3 trials at pH 6.5 and 10.5 plotted with data from blank runs at pH 6.5 and 10.5 for (a) total hypochlorite (OCl− + HOCl) and (b) chlorate ClO3− with an initial NaOCl concentration of 80 mM, 80 °C, and catalyst concentration of 10 μM.
Figure 8. Measured flow rate of oxygen from the IrCl3 trials at (a) pH 6.5 and (b) pH 10.5 with an initial NaOCl concentration of 80 mM, 80 °C, and catalyst concentration of 10 μM.
conclusions about the form of the catalyst can therefore not be made. At first glance, the catalyzed hypochlorite decomposition at pH 10.5 presented in Figure 7a can be believed to be independent of hypochlorite concentration and therefore be of zero order; however, comparison with the oxygen formation during this trial as presented in Figure 8b paints a different picture. It can be seen that the oxygen formation continues long after the depletion of hypochlorite, which would have been at t = ∼35 min if the decomposition would have been a purely zero 3772
DOI: 10.1021/ie504890a Ind. Eng. Chem. Res. 2015, 54, 3767−3774
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Industrial & Engineering Chemistry Research
The addition of 10 μM of FeCl3, Fe3O4, CeCl3, Na2Cr2O7, Na2MoO4, RuCl3, and RuO2, 100 ppm of AgCl, and 9 ppm of Al2O3 to dilute NaOCl solutions did not catalyze the decomposition of hypochlorite to oxygen or chlorate. However, addition of 10 μM of CoCl2 or IrCl3 catalyzed the formation of oxygen. IrCl3 was also found to catalyze the formation of chlorate, especially at the higher pH of 10.5.
order reaction. The lack of hypochlorite in data points above 40 min is more likely due to limitations in the IC detector sensitivity rather than an actual complete depletion of hypochlorite. The catalysis of chlorate at pH 6.5 can be seen during the first few minutes of the reaction (Figure 7b), where the measured chlorate concentrations are significantly higher than those of uncatalyzed chlorate formation (see Table 1 for standard deviation data for the uncatalyzed formation). At both pH 6.5 and 10.5, an increase in oxygen formation can be seen late in the reaction (Figure 8). A similar broad peak was not observed in the oxygen evolution curve for any other experimental conditions but was observed in all experiments involving iridium chloride. This second peak in the oxygen evolution rate can perhaps be explained by considering the decrease in the hypochlorite concentration over time, as the appearance of this peak coincides with the probable switch from a zero order decomposition of hypochlorite to a higher order. Furthermore, the decrease in hypochlorite concentration will result in a change in the oxidative potential of the electrolyte, which could have an effect on the added iridium catalyst. Iridium in aqueous solution at room temperature has an oxidized form at high pH (reaction 19) or oxidizing potential according to the Pourbaix diagram.18 IrO2 + 2H 2O ⇌ IrO4
2−
+
+ 4H + 2e
−
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Information regarding instruments and chemicals used in the study, description of the experimental procedure, method for the pKa determination, method validation, and possible sources of error in the measurements. This material is available free of charge via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS Financial support from the Swedish Energy Agency, the Nobel Research Fund, Permascand AB, and AkzoNobel PPC is gratefully acknowledged. Researchers at Permascand AB and AkzoNobel PPC involved in the present study are acknowledged for useful discussions. Johan Wanngård at AkzoNobel PPC is acknowledged for help with the pKa measurements. Inger Odnevall Wallinder and Xian Zhang at the Department of Surface and Corrosion Science at KTH are acknowledged for letting us use their HPLC instrument and for their help during use. Åsa Björkbacka at the Department of Applied Physical Chemistry at KTH is acknowledged for her help with the ICP measurements.
(19)
2−
If this oxidized IrO4 species, where the Ir oxidation state is 6+, is the only one responsible for catalysis of chlorate formation, the effect of pH on the chlorate catalysis selectivity can be explained. Furthermore, it could also explain the second broad peak in the rate of oxygen evolution. As the concentration of HOCl decreases, the oxidative potential of the solution decreases according to the Nernst equation for HOCl (eq 20),18 E = 1.494 + 0.0295 × pH + 0.0295 × log
ASSOCIATED CONTENT
S Supporting Information *
[HOCl] [Cl−]
■
(20)
so that the equilibrium is driven to the left side of eq 19. The result is an increase in concentration of IrO2, which could be active for the decomposition of hypochlorite to yield oxygen, explaining the appearance of a second peak in detected oxygen. The results suggest that the catalysis of chlorate formation could be exclusive to higher oxidation states of Ir.
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CONCLUSIONS An experimental setup that allows for time-resolved gas and liquid concentration measurements under controlled pH has been designed and applied for the study of catalyzed and uncatalyzed hypochlorite decomposition. The rate of oxygen formation from homogeneous hypochlorite decomposition in the pH range 4−9 has a maximum at pH ∼6.5 and thus depends on the concentrations of both HOCl and OCl−. This pH coincides with the maximum rate for chlorate production, and it is suggested that the same reaction intermediate is involved in the formation of both chlorate and oxygen. Rate expressions for overall hypochlorite decomposition, chlorate formation, and oxygen formation at pH 6.5 and 80 °C were obtained. All three rate expressions were of the form ri = ki[HOCl]2[OCl−],with rate constants ki of 2.39 ± 0.12 M−2 s−1, 0.731 ± 0.035 M−2 s−1, and 0.046 ± 0.001 M−2 s−1, respectively.
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NOMENCLATURE IC = ion chromatography MS = mass spectrometry ICP-AES = inductively coupled plasma atomic emission spectrometry H = total hypochlorite (HOCl + OCl−) Ci = concentration of species i FO2 = molar flow rate of oxygen at MS detector divided by electrolyte volume Ri,ClO3− = initial reaction rate of chlorate ri = reaction rate of species i ki = reaction rate constant of species i REFERENCES
(1) Schlag, S. CEH Product Review: Sodium Chlorate; IHS Chemical, 2012. (2) Vogt, H.; Balej, J.; Bennett, J. E.; Wintzer, P.; Sheikh, S. A.; Gallone, P. Ullmann’s Encyclopedia of Industrial Chemistry; Wiley-VCH Verlag GmbH & Co. KGaA, 2000. (3) Colman, J. E. Tutorial Lectures in Electrochemical Engineering and Technology. AIChE Symp. Ser. 1981, 77, 244. (4) Burney, H. S. Membrane Chlor-Alkali Process. Mod. Aspects Electrochem. 1993, 393−438. (5) 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.
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Industrial & Engineering Chemistry Research (6) Adam, L. C.; Gordon, G. Hypochlorite Ion Decomposition: Effects of Temperature, Ionic Strength, and Chloride Ion. Inorg. Chem. 1999, 38, 1299−1304. (7) Lister, M. W. Oxygen evolution from sodium hypochlorite solutions. Can. J. Chem. 1962, 40, 729−733. (8) 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. (9) Adam, L. C.; Fábián, I.; Suzuki, K.; Gordon, G. Hypochlorous Acid Decomposition in the pH 5−8 Region. Inorg. Chem. 1992, 31, 3534−3541. (10) Chirnoaga, E. CCXXII. The catalytic decomposition of solutions of sodium hypochlorite by finely divided metallic oxides. J. Chem. Soc. 1926, 129, 1693. (11) Lewis, J. R. The Catalytic Decomposition of Sodium Hypochlorite Solutions. I. Mechanism of the Reaction. J. Phys. Chem. 1928, 32, 243−254. (12) Lister, M. W. Decomposition of Sodium Hypochlorite: The Catalyzed Reaction. Can. J. Chem. 1956, 34, 479−488. (13) Beer, H. B. The Invention and Industrial Development of Metal Anodes. J. Electrochem. Soc. 1980, 127, 303C. (14) Cornell, A. Ruthenium Dioxide as Cathode Material for Hydrogen Evolution in Hydroxide and Chlorate Solutions. J. Electrochem. Soc. 1993, 140, 3123. (15) Hummelgård, C.; Karlsson, R. K. B.; Bäckström, J.; Rahman, S. M. H.; Cornell, A.; Eriksson, S.; Olin, H. Physical and electrochemical properties of cobalt doped (Ti,Ru)O2 electrode coatings. J. Mater. Sci. Eng. B 2013, 178, 1515−1522. (16) Gustavsson, J.; Li, G.; Hummelgård, C.; Bäckström, J.; Cornell, A. On the suppression of cathodic hypochlorite reduction by electrolyte additions of molybdate and chromate ions. J. Electrochem. Sci. Eng. 2012, 2, 185−198. (17) Isa, I.; Shibuya, M.; Ebisawa, M. Manufacturing chlorine dioxide with thallium and silver or palladium catalysts. US Patent 4.169.134, 1979. (18) Pourbaix, M. Atlas of Electrochemical Equilibria in Aqueous Solutions; Pergamon Press Ltd., 1966. (19) Palmer, W. Experimental Inorganic Chemistry; Cambridge University Press, 1970. (20) Knibbs, N. V. S.; Palfreeman, H. The theory of electrochemical chlorate and perchlorate formation. Trans. Faraday Soc. 1920, 16, 402. (21) Hammar, L.; Wranglén, G. Cathodic and Anodic Efficiency Losses in Chlorate Electrolysis. Electrochim. Acta 1964, 9, 1−16. (22) 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. (23) Lister, M. W. The Decomposition of Hypochlorous Acid. Can. J. Chem. 1952, 30, 879−889. (24) Wanngård, J. In Modern Chlor-alkali technology; Wellington, T. C., Ed.; Elsevier Applied Science, 1992; Vol. 5; pp 295−306. (25) 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. (26) 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.
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