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Ind. Eng. Chem. Res. 2008, 47, 5284–5290
Reactions and Kinetics of Cl(III) Decomposition Tuula Lehtimaa,*,† Ville Tarvo,‡ Ge´rard Mortha,§ Susanna Kuitunen,‡ and Tapani Vuorinen† Department of Forest Products Technology, TKK Helsinki UniVersity of Technology, P.O. Box 6400, 02015 Espoo, Finland, Department of Biotechnology and Chemical Technology, TKK Helsinki UniVersity of Technology, Finland, and EFPG, France
Several pathways leading to the decomposition of chlorite and chlorous acid have been published. In this study, both experimental and computational approaches have been applied to clarify the authenticity of the different routes. The decomposition of chlorine (III), i.e. chlorous acid and chlorite, was monitored with iodometric titration at changing chloride concentrations, temperature, and existence of iron (III) at pH 1-3. Dimethylsulfoxide (DMSO) was used to prevent hypochlorous acid from reacting with chlorite. Chlorine dioxide was not formed in the absence of metals when hypochlorous acid and chlorine were trapped. The self-decomposition of Cl(III) proceeds only via the acidic form while chlorite is stable. Chloride ions enhanced the chlorous acid decomposition rate especially at low pH. Chlorite decomposes in the presence of Fe3+ ions. General kinetic parameters and their temperature dependencies were determined for chlorous acid selfdecomposition. Introduction Chlorine dioxide (ClO2) is widely used in pulp bleaching and water disinfecting.1–3 In both applications, chlorine dioxide has substituted chlorine, because chlorine dioxide treatment does not produce trihalomethanes (THM) which are carcinogenic byproduct associated with chlorination of pulp or naturally occurring organics in raw water.3–5 However, chlorine dioxide treatment results in formation of some chlorinated organics, both in pulp bleaching and water treatment, but their amounts are lower than with chlorine.1,5 Reactions of chlorine dioxide result in the formation of several inorganic chlorine compounds. These species are reactive with organic compounds and each other. As a consequence of their formation the applications that use chlorine dioxide comprise of a complex mixture of different organic and inorganic reactions. The initial reaction is a fast reaction of chlorine dioxide (ClO2) with phenolic structures.6,7 This reaction produces chlorite (ClO2-) and hypochlorous acid (HOCl) in 1:1 molar ratio.8 Formed chlorite can decompose in acidic conditions and form hypochlorous acid and chlorate (ClO3-).9–13 Chlorite also reacts very rapidly with hypochlorous acid forming an intermediate Cl2O2, which reacts further forming chlorate and chlorine dioxide.10,12–16 Hypochlorous acid is in equilibrium with chlorine (Cl2). Both of these compounds (HOCl and Cl2) chlorinate various organic structures, resulting in chlorinated organic compounds that can be harmful for the environment.17–19 To enable the simulation or prediction of the behavior of all of these compounds a computer model based on real chemistry and kinetics is needed. However, the literature published so far gives very varying information on the kinetics as well as on the reactions included in the decomposition of chlorite. Since the amount of chlorite that is decomposed also affects how HOCl reacts, that is with chlorite or with organic compounds forming chlorinated organics, it is vital to obtain precise data * To whom correspondence should be addressed. E-mail:
[email protected]. † Department of Forest Products Technology, TKK Helsinki University of Technology. ‡ Department of Biotechnology and Chemical Technology, TKK Helsinki University of Technology. § EFPG.
for this reaction. Also the amount of chlorate formed through the decomposition of Cl(III) is crucial because it is known to be toxic to certain types of algae1 and it may possess reproductive toxic effects.20 Thus its amount in waste waters and in drinking water should be minimized. The amount of chlorate formed is important also in pulp bleaching; chlorate is known to be inactive in bleaching conditions,1 thus the more chlorate is formed, the more oxidizing power is lost. In neutral conditions, half of the used chlorine dioxide can end up as chlorite.5 Since also chlorite itself is suspected to be a reproductive toxin20 it is important to have more information of its decomposition rate. In an aqueous solution, chlorite is in equilibrium with its acidic form, chlorous acid (HClO2), and together they form the total amount of chlorine (III).21 Several papers have been published on the reactions of Cl(III) and it is obvious that the numerous reactions are complex and involve many different inorganic chlorine compounds.14,21–24 Only a limited amount of studies have concentrated on the decomposition of Cl(III) in the absence of reactive intermediates that are formed during the decomposition.9,10,12,15 Chlorous acid and its dissociated form, chlorite, are in equilibrium according to eq 1. The Ka value of this equilibrium has been reported to be (at room temperature) 1.1 × 10-2 M.11,21 HClO2 / ClO2- + H+
(1)
Sometimes the decomposition of Cl(III) is referred to as the net reaction given by eqs 2–4.11,21,25,26 However, these reactions producing chlorine dioxide and chlorate are a sum of a series of reactions.9,10,13,27 In addition to Cl(III) decomposition, they include several follow-up reactions of the formed reactive species. The stoichiometries of the subsequent reactions are dependent on reaction conditions, and therefore, eqs 2–4 do not describe the decomposition accurately except in the conditions where they were determined. 5HClO2 f 4ClO2 + Cl- + 2H2O + H+
(2)
5ClO2- + 4H+ f 4ClO2 + Cl- + 2H2O
(3)
3HClO2 f 3H+ + 2ClO3- + Cl-
(4)
10.1021/ie0714089 CCC: $40.75 2008 American Chemical Society Published on Web 06/25/2008
Ind. Eng. Chem. Res., Vol. 47, No. 15, 2008 5285
Several reaction routes (eqs 5–15) have been presented for the decomposition of chlorous acid without the participation of intermediate components. Reaction 5 has been suggested to occur in the absence of chloride ions.9–11,13 2HClO2 f HOCl + H+ + ClO3-
(5)
A reaction between chlorous acid and chlorite has been under discussion.10,11 Hong and Rapson9 suggested that the reaction proceeds according to eq 6. HClO2 + ClO2- f HOCl + ClO3-
(6)
The catalyzing effect of chloride ions in Cl(III) decomposition has been reported by many authors.9,10,12,13 Chloride has been claimed to react with chlorous acid according to reaction 7,9,10,12,13 while Schmitz and Rooze15 suggested that the reaction proceeds in a stepwise manner (eqs 8 and 9). HClO2 + Cl- + H+ f 2HOCl
(7)
HClO2 + Cl- + H+ f Cl2O + H2O
(8)
Cl2O + H2O f 2HOCl
(9)
27
Kieffer and Gordon presented an alternative reaction for the chloride-catalyzed decomposition. They assumed that protons are not involved in the reaction and that the reaction proceeds through an intermediate structure (eqs 10 and 11). HClO2 + Cl- / [HCl2O2-] -
-
[HCl2O2 ] + Cl
(10)
f products (rate-determining)
(11)
Horvath et al.12 proposed new routes for the Cl(III) decomposition, proceeding via a dichlorotrioxide (Cl2O3) intermediate, eqs 12–15. Reactions 12 and 13 are kinetically and by stoichiometry equivalent to reaction 5. In addition, Cl2O3 reacts with chloride and chlorous acid to yield various chlorine compounds (eqs 14 and 15). 2HClO2 / Cl2O3 + H2O +
(12) -
Cl2O3 + H2O f H + HOCl + ClO3 -
+
(13)
Cl2O3 + Cl + H f Cl2O2 + HOCl
(14)
Cl2O3 + HClO2 + H2O f 3H+ + Cl- + 2ClO3-
(15)
Fe3+ ions have been reported to catalyze the decomposition of Cl(III) even at very low concentrations.28–35 Since unbleached kraft pulp typically contains metals such as Mn, Fe, and Cu,36 this effect can have some significance in the first chlorine dioxide bleaching stage. The rate of the Fe3+ ion catalyzed reaction increases as the pH increases from 0.3 to 2.7.28,30–37 The effect of iron is not seen above this pH range because at higher pH the Fe3+ ions start to hydroxylate and form polynuclear compounds.28,38 As shown above, several reactions have been suggested in literature for the decomposition of chlorite and chlorous acid. This study was performed with the intention to clarify the reactions involved in the decomposition of Cl(III) without the effect of intermediately formed species. The aim was also to define general kinetic parameters for those reactions and their temperature dependencies. Previous studies have reported kinetic parameters at 25 °C,9,12,15,27 but information of kinetic constants at elevated temperatures is lacking. This work covers elevated temperatures and applies a nonlinear curve fitting technique in the kinetic analysis. Using a simultaneous curve fitting method, instead of conventional pseudo-first-order techniques, brings
Table 1. Ka,1 Values Used in Parameter Fitting T/°C Ka,1/M
20 0.011
40 0.0078
50 0.0065
60 0.0056
70 0.0048
additional credibility for the determined reaction parameters.39 Dimethylsulfoxide (DMSO) was used to scavenge the reactions of chlorous acid with the formed hypochlorous acid or chlorine. This scavenger is known to trap hypochlorous acid and chlorine effectively, but it does not affect other chlorine species.40–42 Experimental Section Reagents in the Reactor. Fresh sodium chlorite solution (NaClO2, Fluka, ∼80%) was prepared daily. Other reagents used were sodium chloride (NaCl, Acros Organics, min 99.5% p.a.), perchloric acid (HClO4, Prolabo, min 70-72%), dimethylsulfoxide (DMSO, Chimie Plus Laboratoires, min 99%), iron sulfate (Fe2(SO4)3 · xH2O, Aldrich, min 97.0%), copper sulfate (CuSO4 · 5H2O, Prolabo, min 99%), and EDTA (Acros Organics, min 99%). All chemicals were used as obtained. Reagents for Titration. The titration stock solution (0.1 M) was prepared from sodium thiosulfate (Na2S2O3 · 5H2O, Prolabo, min 99.5%), and its concentration was verified weekly by titrating with potassium iodate (KIO3 Merck, min 99.7%). The dilution to 0.01 M solution was done daily. The borate buffer used to adjust the pH was prepared from ortho-boric acid (H3BO3, Prolabo, ∼99.5%) and sodium hydroxide (NaOH Carlo Erba, min 97%). Also potassium iodide (KI, Roth, min 99%) and hydrochloric acid (HCl, Aldrich, 37 wt %) were used. All chemicals were used as obtained. The used commercial sodium chlorite (∼80%) contains a notable amount of NaCl as an impurity. Its amount was determined with an ion chromatograph (Dionex ICS-1500) and was found to be 18.4%. This was taken into consideration when interpreting the results and in the calculation of kinetic values. The amounts of iron, manganese, and zinc in sodium chlorite were analyzed with ICP-OES equipment (Varian, Liberty). They were found to be Fe 25.7 µg/g, Mn 0.8 µg/g, and Zn 4.3 µg/g. The possible catalyzing effect of these metals was inhibited by the use of EDTA. The Cl(III) decomposition experiments were carried out in a double-wall, thermostatted 200 mL glass reactor equipped with a Teflon lid. A magnetic stirrer was used for mixing. The temperature in the experiments was 60 °C and pH 2 unless otherwise mentioned. Since some carboxylic acids can be reactive with chlorite, buffers were not used to adjust the pH.43,44 For example, it was noticed in our preliminary tests that citrate buffer reacts with chlorite. The pH in the reactor was verified with random pH measurements. It was found that the pH stayed practically constant (
