Stability of chlorine dioxide in aqueous solution

Water Res. Vol. 16, pp. 1379to 1382, 1982 Printed in Great Britain.All rights reserved

0043-1354/82/091379-04503.00/0 Copyright • 1982PergamonPress Ltd

STABILITY OF CHLORINE DIOXIDE IN AQUEOUS SOLUTION MAGDA MEDIR and FRANCESCGIRALT Departaments de Quimica Fisica i Quimica T6cnica, Facultat de Quimica, Universitat de Barcelona, Tarragona, Catalunya, Spain (Received July 1981)

Abstract--Chlorine dioxide in unbuffered aqueous solution at pH 9 and 25°C decomposes to chlorate, chlorite, chloride and oxygen by coupled slow and rapid reactions. The initiation of the second rapid decomposition depends on the initial chlorine dioxide concentration and ionic strength of the solution. Chloride ion has a catalytic and inhibiting effect, changing the product distribution to a 1:1 molar ratio of chlorate and chlorite, thus increasing the oxidizing potential of chlorine dioxide. In buffered solutions at pH 8.8 the reaction is pseudo-second order, yielding chlorate and chlorite as products.

INTRODUCTION The use of chlorine dioxide in waste-water treatment is recommended when dealing with industrial effluents contaminated with phenolic and chlorophenolic compounds which impart undesirable tastes and odors to the water. Chlorine dioxide reacts completely to oxidize the taste-producing compounds many times faster than it does free available chlorine and its efficiency is not impaired by a high pH environment (White, 1972). The stability of aqueous solutions of chlorine dioxide under the conditions prevailing during wastewater treatment processes, i.e. neutral or slightly alkaline unbuffered medium, as well as the reactions undergone by chlorine dioxide under these conditions are not well established and a better understanding is desirable to interpret its role as oxidizing agent and disinfectant. An extensive review on the stability of aqueous solutions of chlorine dioxide has been conducted by Gordon (1972). Chlorine dioxide disproportionates into chlorite and chlorate according to 2C102 + H 2 0 = C10; + C102 + 2H ÷

(1)

in strong alkaline medium, and above pH 6 in the presence of carbonate, phosphate and borate buffers. This reaction is base catalysed, the rate of reaction increasing with pH (Granstom & Lee, 1957; Zolotukhim et al., 1965). In acid medium chloride, in addition to chlorate and chlorite, is also a product of the decomposition of chlorine dioxide (Brown, 1952). The present work studies the decomposition of chlorine dioxide in aqueous unbuffered medium at pH 9 and 25°C. Under these conditions hypochlorous acid is about 97% dissociated and chlorous acid is completely dissociated (Fialdes, 1972), avoiding therefore its own decomposition (Gordon, 1972). The effects of chlorine dioxide initial concentration (0.002-0.04 mol 1-t), ionic strength of the solution by

addition of 0.14 and 0.27 mol 1- t of sodium sulphate and of the presence of sodium chloride (0.1-2.0reel 1- t ) have been studied. Results for the decomposition of chlorine dioxide in a buffered solution are also reported. This paper presents in brief results pertaining to work still in progress because of their practical interest. MATERIALS AND METHODS Reagents

Double distilled water was used for all the solutions. Chlorine dioxide was prepared from 20% sodium chlorite (commercial grade) solution by dropwise addition of 1 N H,SO+. The generated chlorine dioxide was swept from the solution using nitrogen as carrier, passed through sodium chlorite solution to remove any chlorine formed, and dissolved in double distilled water or in aqueous solutions of Na2SO+ and NaCI placed in the reaction vessel. All chemicals were reagent grade and standardized according to standard methods (White, 1972). Procedure and apparatus The reactions were carried out at 25°C in a jacketed pyrex glass reaction vessel of 500 ml protected from light and in the absence of gaseous phase. The pH was maintained constant at 9 by automatic addition of sodium hydroxide from a burette controlled by a pH-stat. The initial concentration of chlorine dioxide was determined by acidneural iodometric titration with sodium thiosulphate and the chlorine concentration in the solutions was negligible. The chlorine dioxide consumption was determined by recording the addition of sodium hydroxide, which is directly related to the chlorine dioxide reacted, as one equivalent of hydrogen ion is formed from each mole of chlorine dioxide decomposed. The dilution was less than 0.4~ (v/v) since less than 2 ml NaOH solution were added. The chlorine dioxide solution was allowed to react and when the reaction was practically finished, as shown by the decreased rate of addition of sodium hydroxide, the chlorine dioxide and product concentrations were determined. The chlorate and chlorite concentrations were determined by iodometric titration with sodium thiosulphate at different pH values (Hong, 1966; Medir, 1973). The chloride concentration was determined by titration with silver nitrate solution using both the Mohr method and also the potentio-

1379

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Time,t(min) Fig. 1. Degree of conversion of chlorine dioxide in water and aqueous sodium sulphate solutions

metric method for detection of end point, after flushing the remaining chlorine dioxide out of the solution. An oxygen probe and a YS1 model 54 oxygen meter were used for the determination of oxygen in very dilute chlorine dioxide solutions. RESULTS AND DISCUSSION

Decomposition in water and aqueous sodium sulphate solutions The rate of decomposition of chlorine dioxide was studied in a series of experiments carried out in water and in 0.14 and 0.27mol1-1 sodium sulphate solutions with initial chlorine dioxide concentrations ranging from 0.002 to 0.04 mol I-1. A plot of the degree of conversion of chlorine dioxide vs time shows in Fig. 1 that there is an initial slow reaction, followed by a more rapid one. This behaviour was also observed by Heijne & Teder (1973) in experiments carried out at pH 2-7 and 40-80°C, in unbuffered solutions containing 0.0014).01 mol 1-1 of chlorine dioxide. However, this reaction path seems to be different from that reported by G o r d o n (1972) in strong alkaline medium where the initial slow decomposition seemed not to be present. The products of the decomposition at pH 9 have been found to be chlorate, chlorite, chloride and oxygen. The average relative amounts of these products for 9 mol of reacting chlorine dioxide are 5, 3, 1 and 3 respectively, with a maximum relative error of 4~o. This product distribution obtained in water and in sodium sulphate solutions does not correspond to the typical disproportionation reaction (1), which takes place at high alkaline concentrations or in the presence of buffers. The rate of decomposition as well as the product distribution show the complexity of the reacting system. The trend of the data suggests a first slow reaction, on which is later superimposed another reaction of a much higher rate. Figure 1 shows that the conversion data pertaining to the first slow reaction

present an inflection point where the second reaction starts accelerating with an autoeatalytic shape. In order to deduce a kinetic model for this complex reacting system it is necessary to analyze the intermediate product distribution during the course of the reaction and research is still underway. However, some information may be deduced from the present data, which is of great practical interest. The initial rate of the slow reaction plotted in Fig. 2 vs the initial concentration of chlorine dioxide, is of pseudo-first order, ro =

d[CIO2] ,=o = k°[CIO2]° dt

(2)

with a rate constant k ° = 0.3 x 10-3 min-~ for the decomposition reaction in water, while for aqueous sodium sulphate solutions of ionic strength 0.65-0.90 mol l - 1 is k ° = 0.87 × 10- 3 m i n - 1. The standard deviation is s = 0.08 x 10 -3 for each rate constant. The rate of decomposition of chlorine dioxide is more reproducible under conditions of constant ionic strength.

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Stability of chlorine dioxide in aqueous solution

characteristic disproportionation reaction (1) of chlorine dioxide in strong alkaline medium. The presence of sodium chloride changes the product distribution decreasing the formation of chlorate and therefore enhancing the oxidizing potential of chlorine dioxide. Rapson & Anderson (1977) observed that the addition of sodium chloride increased the efficiency of chlorine dioxide bleaching of pulp, yielding higher brightness and decreasing chlorate formation significantly, which agrees with the present behaviour of chlorine dioxide. Aqueous solutions of chlorine dioxide containing sodium chloride are less stable. The time of initiation of the second reaction may be correlated to the initial concentration of chlorine dioxide by

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Fig. 3. Degree of conversion of chlorine dioxide in aqueous sodium chloride solutions.

The time of initiation of the second fast reaction t (min), which causes most of the decomposition of chlorine dioxide, decreases with the chlorine dioxide initial concentration [C102]0 (mol 1-t) according to t = a [C102]o 0"72

1381

(3)

where a = 4.0 for the decomposition in water and a = 2.4 for the decomposition in aqueous sodium sulphate solutions. These results show that aqueous solutions of chlorine dioxide at pH 9 and 25°C are fairly stable for a certain period of time, which depends on the initial concentration of chlorine dioxide. At high initial concentrations the second rapid decomposition takes place sooner and at a faster rate than at low concentrations. These initial concentration effects are even more noticeable with the addition of inert salts, such as sodium sulphate. Under these conditions chlorine dioxide solutions are less stable.

Decomposition in aqueous sodium chloride solutions Chlorine dioxide may be used in waste-water treatment processes where the inflowing waters contain sodium chloride. The study of the chlorine dioxide decomposition in the presence of sodium chloride has a practical as well as a theoretical interest since chloride ion is a product of the reaction in pure aqueous solutions. Experiments were carried out in sodium chloride concentrations 0.1, 0.4, 0.8 and 2.0 mol 1-1 with initial chlorine dioxide concentrations ranging from 0.002 to 0.02 mol 1-1, i.e. with sodium chloride in excess. In the presence of sodium chloride the decomposition reaction follows a different path, as shown in Fig. 3 where the degree of conversion of chlorine dioxide is plotted vs time. There is a slow reaction of a few minutes duration, which is not appreciable on this graph, followed by a second faster reaction, which does not present an autocatalytic shape. The products of the reaction are chlorate and chlorite in a 1 : 1 molar ratio, which coincides with the

t = 0.0013 [C102] o 1.66

(4)

being up to 13rain for [ C l O 2 ] o = 0 . 3 x 10 -2 tool l - 1. Neglecting the first slow reaction period the rate of decomposition is of pseudo-first order with a rate constant k = 3.9 x 10- 3 min- 1 for all NaCl solutions of ionic strength 0.1-0.8 mol l-1. The conversion data (f) are well fitted up to 200 min with SD = 0.02. This is a simple model which may be used as a first approximation for design purposes. Comparison of Figs 1 and 3 shows that the second reaction occurs earlier but at a lower rate when sodium chloride is present. Chloride ion seems to have an inducing and inhibiting effect on the overall decomposition reaction. Different decomposition rates were observed when initial chlorine dioxide and sodium chloride concentrations were comparable. In these cases the triggering time for the second reaction is practically independent of the initial chlorine dioxide concentration and, therefore, cannot be predicted by equation (4). However, sodium chloride concentrations 10-2 M are sufficient to yield a second reaction behaviour of the same type as that reported for high concentrations in Fig. 3.

Decomposition in buffered solution Chlorine dioxide at 25°C in a buffered sodium bicarbonate solution at pH8.8 disproportionates according to equation (1). The rate of decomposition is pseudo-second order with a rate constant k = 4.751 mol- t min- 1. The rate of chlorite formation is one half that of chlorine dioxide decomposition. A second order rate expression was also reported by Zolotukhim et al. (1965) for the chlorine dioxide decomposition in phosphate buffer. CONCLUSIONS

Aqueous solutions of chlorine dioxide are fairly stable at 25°C and pH 9 for an initial period of time before a fast decomposition takes place. The length of the initial stable period decreases with increasing chlorine dioxide concentration and in the presence of

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MAGDA MEDIR and FRANCESCGIRALT

inert electrolytes. The reaction products are chlorate, chlorite, chloride and oxygen which are different from those previously reported in strong alkaline medium. Addition of sodium chloride reduces significantly the induction time but slows down the second reaction and changes the product distribution to the one obtained in strong alkaline medium, where equal amounts of chlorate and chlorite are formed. In buffered solution at pH 8.8 the reaction is pseudo-second order with chlorate and chlorite as products. Acknowledoements--One of the authors M. Medir would like to thank Professor W. H. Rapson, University of Toronto, for his help and guidance during the initial part of this work.

REFERENCES Brown R. W. (1952) Thermal decomposition of chlorine dioxide. Tappi 35, 75-80.

Fialdes E. (1972) Dechlorination des eaux sur carbon actif. Inform. Chim. 100, 127-131. Gordon G., Kieffer R. G. & Rosenblatt D. M. (1972) The chemistry of chlorine dioxide. Progress in Inorganic Chemistry (Edited by Lippard S. J.), Vol. 15, pp. 201-287. Wiley, New York. Granstrom M. L. & Lee G. F. (1957) Rates and mechanisms of reactions involving oxychloro compounds. Publ. Wks. 88, 90-92. Heijne von G. & Teder A. (1973) Kinetics of the decomposition of aqueous chlorine dioxide solutions. Acta chem. scand. 27, 4018-4019. Hong C. C. (1966) Reactions of chlorous acid. Ph.D. Thesis, University of Toronto. Medir M. (1973) Reactions of chlorine dioxide with water at pH 9. M.A.Se., University of Toronto. Rapson W. H. & Anderson C. B. (1977) Improving the efficiency of chlorine dioxide bleaching. Pulp Paper Can. 3, 52-55. White G. C. (1972) Handbook of Chlorination. Van Nostrand Reinhold, New York. Zolotukhim V. M., Flis I.E. & Mischenko K. P. (1965) Electrochemical studies of the kinetics of the decomposition of chlorine dioxide in aqueous solution. Khim. i Khim. Technol. 8, 764-767.