Ind. Eng. Chem. Res. 1997, 36, 11-16
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Influence of Mercuric Nitrate on Species and Reactions Related to Chlorine Dioxide Formation Bhart Indu, M. Fazlul Hoq, and W. R. Ernst* School of Chemical Engineering, Georgia Institute of Technology, Atlanta, Georgia 30332
Mercuric ions influence reactions and intermediates that are involved in forming chlorine dioxide from chlorate ions. Addition of mercuric ions to reaction solutions can aid in understanding the mechanism and kinetics of this system. Mercuric ions do not react with aqueous solutions of chlorine dioxide unless those solutions contain chlorous acid. This unusual effect has enabled us to confirm that chlorous acid is an intermediate in the formation of chlorine dioxide in the methanol-chlorate reaction. This work discusses the effect of mercuric ions on solutions containing various Cl species (chlorine, chlorine dioxide, chlorous acid, and both chlorine dioxide and chlorous acid) and on reactions that involve these same species. The work also explores the methanol-chlorate reaction in the absence of mercuric ions. In the initial stage of the process, chlorine dioxide forms after an induction period, during which an intermediate rapidly forms, maximizes in concentration, and then disappears, and a second intermediate forms and attains a steady-state concentration. We show by several methods, including mercuric ion addition, that the second intermediate is chlorous acid. We have not conclusively identified the first intermediate. The increase and decline in its concentration is accompanied by the development and disappearance of the solution color (yellow). Although the UV absorbance behavior of the system suggests that the intermediate is chlorine, the color of the solution associated with this intermediate is not characteristic of solutions containing chlorine at these concentrations. Introduction Because of environmental concerns, chlorine dioxide is replacing chlorine in many chemical oxidation applications. Two of the most important applications are bleaching of pulp in the manufacture of paper and water treatment. Although there are several ways of synthesizing chlorine dioxide, most commercial operations react chlorate ions in acidic solution with a reducing agent such as methanol, hydrogen peroxide, sulfur dioxide, or sodium chloride. For smaller applications, chlorine dioxide can be produced by oxidizing sodium chlorite with chlorine (Masschelein, 1979). Lenzi and Rapson (1962) proposed that chlorate undergoes sequential reduction when reacted with a reducing agent, R:
mechanism. Reaction products, chlorine dioxide and chlorine, are routinely monitored in reaction studies by ultraviolet spectroscopy (see Hong et al., 1967; Burke et al., 1993; Tenney et al., 1990). Recently, Ernst et al. (1996) reported results of a kinetic study of what we suspect is the initial step of the reduction of chlorate by methanol in acidic solution:
H+ + CH3OH + ClO3- f HClO2 + HCHO + H2O (7)
HClO3 + R f HClO2 + RO
(1)
HClO2 + R f HClO + RO
(2)
HClO + R f HCl + RO
(3)
The reaction was continuously monitored in an NMR apparatus. To suppress the influence of side reactions, we added mercuric nitrate to the reaction solutions. We had previously recognized the potential use of mercuric ions as a tool for studying reactions in the production of chlorine dioxide; however, we have not previously provided details of this method. This paper summarizes our observations of the influence of mercuric ion on a variety of species and reactions that are related to the chemistry of chlorine dioxide production.
HClO3 + HCl f HClO2 + HClO
(4)
Background
HClO3 + HClO2 f 2ClO2 + H2O
(5)
HCl + HClO T Cl2 + H2O
(6)
According to this mechanism, chlorine dioxide is produced by the reaction of the intermediate, chlorous acid, with chlorate ions in step 5. The reaction is usually run at high acidity, at which the stable form of chlorine is Cl2 (step 6). The intermediates, chlorous acid and hypochlorous acid, are highly reactive, making it difficult to experimentally investigate rates of the individual steps of the * Author to whom correpondence should be addressed. Phone: 404-894-2878. Fax: 404-894-2866. S0888-5885(96)00431-9 CCC: $14.00
Lenzi and Rapson (1962) observed that, under suitable conditions of acidity and mixing, chlorine dioxide is the predominent product of the reaction of sodium chlorate and sodium chlorite in solution. When they incorporated silver sulfate into an experiment under identical mixing conditions, chlorine was produced instead of chlorine dioxide. The latter experiment was conducted to illustrate the importance of chloride ions as an intermediate in the production of chlorine dioxide. That work motivated us to explore other chemicals that would influence intermediates in chlorine dioxideproducing reactions. A negative aspect of silver ion addition for use in kinetic experiments is that the product silver chloride is extremely insoluble. Reacting solutions become cloudy as a precipitate forms, making it difficult to monitor formation of chlorine dioxide and © 1997 American Chemical Society
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Ind. Eng. Chem. Res., Vol. 36, No. 1, 1997
chlorine by the usual continuous ultraviolet spectroscopic methods. Mercuric ions, Hg2+, remove (sequester) chloride ions from solution; however, unlike silver ion, mercuric ions form a soluble product, mercuric chloride, HgCl2. The solubility of mercuric chloride in aqueous solution is 0.269 mol/L of water (or 73 g/L of water) as reported by Coetzee and Ritchie (1969). In aqueous solution, mercuric chloride exists exclusively (ca. 99%) as HgCl2 molecules, although some hydrolysis occurs as
HgCl2 + H2O S Hg(OH)Cl + H + Cl +
-
(8)
with an equilibrium constant, log(Keq) ) -9.56 ( 0.05 (Ciavatta and Grimaldi, 1968). The soluble mercuric chlorides do not interfere with the UV absorbance of chlorine dioxide in solution. Therefore, mercuric ions can be conveniently incorporated in batch kinetics experiments monitored by UV spectroscopy. Experimental Section Certified reagents, crystalline sodium chlorate, concentrated sulfuric acid, formic acid (88%, 23.6 M), formaldehyde (37% w/w, 10-15% v/v methanol), mercuric chloride, anhydrous sodium sulfate, and nitric acid were purchased from Fisher Scientific, Inc. Sodium chlorite was purified by fractional crystallization of technical-grade material (Kodak, Inc.). The third fraction of crystals was used in this work. No chlorate was found in the crystals using a titration procedure of Aieta et al. (1984). A stock solution of chlorine dioxide was prepared by reaction of oxalic acid with chlorate in a sulfuric acid solution (Masschelein, 1979). Before mercuric nitrate (ACS reagent, Aldrich) was added to any reaction solution, it was first dissolved in a small amount of nitric acid. One or more of the following instruments or procedures was used to monitor the progress of reactions in the various studies: (1) Milton Roy Spectronic 1201 UV spectrophotometer with either a static cell or flow cell (both with a 1 cm path length) to monitor chlorine dioxide at 357 or 370 nm, chlorine at 322 nm, and chlorous acid at 250 nm. The equipment has been described in previous studies (Hoq et al., 1991; Indu et al., 1991; Burke et al., 1993). (2) Varian XL-400 Fourier transform proton NMR spectrometry for analysis of organic reagents or products. This instrument and procedures are described in Ernst et al. (1996). (3) Potentiometric titration to measure chlorine dioxide, chlorine, chlorite, and chlorate in water (Aeita et al., 1984). The method involves a potentiometric determination of iodine formed by the oxidation of iodide by the various chlorine-containing species. The titration was conducted at two different pH’s and with and without sparging of nitrogen through the solution. The method was tested by titrating a prepared mixture of the four chlorine-containing species of known concentration. (4) Visual observation. In a few experiments, a color change was sufficient indication of reaction. In some experiments, the UV spectrophotometer with flow cell was incorporated into a batch recirculation reactor apparatus. The reaction vessel was a 125-mL stirred flask, partially submerged in a constant-temperature bath ((0.5 °C). The total volume of the reaction solution was 100 mL in each experiment. Liquid was continuously pumped at 100 mL/min from the flask through the flow cell and back to the flask. The residence time in this recycle system was about 6
s. To start this procedure, a solution of all reagents except one was placed in the flask, and the stirring, flow through the UV cell, and monitoring of UV absorbance at appropriate wavelengths were begun. At time zero, the final ingredient was added to the flask and the absorbance-time behavior of the system continuously monitored. (Also see Burke et al., 1993.) Results and Discussion 1. Effect of Mercuric Nitrate on Various Chlorine-Containing Species in Solution. a. Chlorine Solution. A 0.02 M solution of chlorine in 2 M sulfuric acid was prepared by bubbling gaseous chlorine through the solution from a cylinder. The solution exhibited a color that is characteristic of chlorine and a maximum UV absorbance at 322 nm. Mercuric nitrate was added at a concentration of 0.02 M. Immediately we observed a sudden disappearance of color and a sharp decline in absorbance at 322 nm. This result is not surprising in that the procedure is similar to that used in preparing pure solutions of hypochlorous acid (see Adam et al., 1992). The reactions involved are well-known:
Cl2 + H2O f HOCl + H+ + Cl-
(9)
Hg2+ + 2Cl- f HgCl2
(10)
b. Acidic Solutions of Chlorine Dioxide, Chlorine Dioxide and Methanol, and Chlorine Dioxide and Chlorate. The following solutions were prepared by combining various reagents and aliquots of a freshly prepared chlorine dioxide stock solution and distilled water for dilution: (1) 0.002 M chlorine dioxide in 1 M sulfuric acid; (2) 0.0015 M chlorine dioxide and 0.24 M methanol in 3 M sulfuric acid; (3) 0.0015 M chlorine dioxide and 3 M sodium chlorate in 3 M sulfuric acid. These solutions each exhibited a maximum UV absorbance at 357 nm typical of chlorine dioxide. A mercuric nitrate solution was added to these respective solutions at the following concentrations: (1) 0.02 M; (2) 0.1 M; (3) 0.1 M. After this addition, the color of the solution and the absorbance at 357 nm did not change significantly, except to reflect a dilution effect from the mercuric nitrate solution addition. This result shows that these chlorine dioxide solutions in the presence of mercuric nitrate are stable. c. Chlorous Acid Solutions. Four solutions containing 0.0044 M chlorous acid and 1 M sulfuric acid were prepared. Mercuric nitrate was added to three of the solutions to a concentration of 0.04 M, and no mercuric nitrate was added to the fourth. All solutions were allowed to react for 5 min. In the solution to which no mercuric nitrate was added, about 8% of the chlorous acid disproportionated to chlorate and chlorine. In the three solutions that contained mercuric nitrate, 95-99% of the chlorous acid reacted to chlorate and chlorine. Although mercuric ions sequester chloride ions, this effect would not interfere in mechanisms of chlorous acid disproportionation in a way that would increase the rate. Two mechanisms of chlorous acid disproportionation have been proposed by Kieffer and Gordon (1968). They referred to one mechanism as an “uncatalyzed” path. The initial, and rate-determining, step of this path is
2HClO2 f HOCl + H+ + ClO3-
(11)
Ind. Eng. Chem. Res., Vol. 36, No. 1, 1997 13
Since chloride is not involved in this rate-determining step, the sequestering of chloride should not influence the rate. Kieffer and Gordon (1968) referred to the other mechanism as the “chloride ion catalyzed” path. The first step is the equilibrium
HClO2 + Cl- T [HCl2O2-]
(12)
which is followed by the rate-determining step
[HCl2O2-] + Cl- f products
(13)
Although they do not discuss the details of reaction 13, it is clear that sequestering of chloride should suppress, rather than enhance, the overall rate of reaction of chlorous acid, because it would prevent the formation of the intermediate in step 12 and eliminate both reactants in step 13. If sequestering of chloride can be ruled out as an explanation for the enhanced rate of chlorous acid decomposition, one must consider alternative effects of mercuric ions in chlorous acid solutions. Previous work has described the influence of ferric ions in chlorous acid (chlorite) solutions. Schmitz and Rooze (1984) have shown that ferric ions catalyze the decomposition of chlorite. On the basis of stopped flow-rapid scan kinetics experiments, Fa´bia´n and Gordon (1992) proposed that catalytic decomposition of chlorite is initiated by formation of FeClO22+ and that the rate-determining step is the redox decomposition of this species. Although there is no previous evidence of a mercuric ionchlorite complex, one might speculate by analogy to the ferric system a mechanism involving rapid formation and decomposition of a complex of the form, HgClO2+. d. Acidic Chlorine Dioxide-Chlorous Acid Solutions. A solution was prepared that contained 0.0015 M chlorine dioxide and 0.05 M mercuric nitrate in 3 M sulfuric acid. This solution had a yellow color and exhibited a maximum UV absorbance at 357 nm. Sodium chlorite was added to the solution to a concentration of 0.022 M. Immediately (in less than a second), the solution color and UV signal at 357 nm disappeared, indicating that all of the chlorine dioxide had reacted. This result and results of the previous two sections indicated that chlorine dioxide did not decompose in the presence of mercuric ions unless chlorous acid was present. This result would suggest that there may be an interaction between chlorine dioxide and chlorous acid, the product of which is susceptible to decomposition in the presence of mercuric ions, or that chlorine dioxide is attacked by products or intermediates formed by the mercuric ion catalyzed decomposition of chlorous acid. Gordon and Emmenegger (1966) proposed a complex, Cl2O4-, in solutions of chlorine dioxide and chlorite ions. They reported that the complex had a dark mahogany color (but which is not evident at the concentrations of our experiments) and formed by the reaction
ClO2 + ClO2- f Cl2O4-
(14)
To explain our observations in terms of this complex, we would have to speculate that the complex, or protonated form, is produced under the highly acidic conditions of our experiments and that mercuric ions either react with it or catalyze its decomposition.
Figure 1. Influence of mercuric nitrate on the chloride-chloratesulfuric acid reaction: Concentration-time profiles for chlorine dioxide (circles) and chlorine (triangles) in two batch reaction experiments at 4.4 M H2SO4, 0.3 M NaClO3, 0.02 M NaCl, and 25 °C. (Open symbols: 0.04 M mercuric nitrate initially added. Solid symbols: without mercuric nitrate.)
2. Effect of Mercuric Nitrate on Various Reactions. a. Reduction of Chlorate Ions by Chloride Ions in Acidic Solution. For these and other reaction experiments, the batch recirculation reactor system, described earlier, was used. Figure 1 shows the concentration-time profiles of chlorine dioxide and chlorine produced in a reaction mixture consisting of 0.3 M sodium chlorate and 0.02 M sodium chloride in 4.4 M sulfuric acid at 25 °C. In the absence of mercuric nitrate, there was a fast reaction producing both chlorine dioxide and chlorine, as shown by the upper two curves. The kinetics and proposed mechanism of this system have been discussed earlier by Hong et al. (1967). In the presence of 0.04 M mercuric nitrate, neither product formed, as shown by the lower two curves. This observation is not surprising since the first step in the proposed mechanism of Hong et al. involves chloride ions. In these experiments, it is evident that the reaction was suppressed because free chloride ions were removed from solution by reaction with mercuric ions as shown by eq 10. A similar experiment was conducted employing a reaction mixture consisting of 0.3 M sodium chlorate and 0.01 M mercuric chloride in 4.4 M sulfuric acid. Over the time period of the first experiment, there was no indication of chlorine dioxide or chlorine production. This result shows that mercuric chloride does not reduce chlorate and that mercuric ions are effective in sequestering chloride ions in this reaction system. Mercuric ions are also effective in sequestering bromide ions. The above procedure was repeated using acidic solutions of sodium bromide and sodium chlorate. The system produced chlorine dioxide when mercuric nitrate was absent. The production of chlorine dioxide was totally suppressed when a mercuric nitrate solution was added. b. Reduction of Chlorate Ions by Methanol in Acidic Solution. (1) Experiments without Mercuric Nitrate. Indu (1993) exploited the influence of mercuric ions in the study of the methanol-chloratesulfuric acid reaction system. Preliminary work that preceded these studies explored the time-dependent behavior of UV absorbance of the methanol-chloratesulfuric acid system in a batch recirculation reactor at 357 nm where the absorbance of chlorine dioxide is a maximum. In a typical experiment the absorbance
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Ind. Eng. Chem. Res., Vol. 36, No. 1, 1997 Table 1. Steady-State Concentration of Chlorous Acid Developed in the Methanol-Chlorate Reaction at 25 °C [HClO2] by
Figure 2. Concentration-time profiles in a batch reaction experiment at 3.3 M H2SO4, 0.4 M NaClO3, 0.23 M CH3OH, and 25 °C for chlorine dioxide (open circles), “chlorine” (closed circles), and chlorous acid (open triangles).
initially increased rapidly, passed through a maximum, declined rapidly to almost zero, and then increased at a gradual rate. During the initial increase and decline of the signal, the solution color in many experiments turned increasingly yellow with an increase in absorbance and became colorless as the signal disappeared. The color returned later when the absorbance gradually increased. Because chlorine also absorbs at 357 nm, we initially assumed that the signal that passed through a maximum represented chlorine. On the basis of the steps of the Lenzi-Rapson model (1962), eqs 1-6, we also assumed that chlorous acid formed during the initial process. On the basis of this reasoning, in subsequent experiments we monitored UV signals at wavelengths 250, 322, and 357 nm, which correspond to characteristic absorbances of chlorous acid, chlorine, and chlorine dioxide, respectively. Each species absorbs at the three wavelengths; therefore, in order to compute concentrations, the signals were uncoupled by solving the three Beer’s law expressions 357 357 357 A357 ) (ClO2 CClO2 + Cl2 CCl2 + HClO2 CHClO2)l
(15)
322 322 322 CClO2 + Cl2 CCl2 + HClO2 CHClO2)l A322 ) (ClO2
(16)
250 250 250 CClO2 + Cl2 CCl2 + HClO2 CHClO2)l A250 ) (ClO2
(17)
The extinction coefficients (M-1 cm-1) corresponding to wavelengths 357, 322, and 250 nm, respectively, are 1171, 726, and 100 for chlorine dioxide; 34, 75, and 35 for chlorine; and 10, 40, and 120 for chlorous acid. In all experiments, a flow cell with a 1 cm path length was used. Figure 2 shows the concentration-time profiles for chlorine dioxide and the other two assumed species, calculated from absorbance measurements using Beer’s law. The graph shows both of the assumed species, chlorine and chlorous acid, formed during an induction period prior to chlorine dioxide formation. “Chlorine” increased rapidly, passed through a maximum, and declined to zero (within the accuracy of the measurements), and chlorous acid increased and rapidly reached a steady-state concentration. The steady-state concentration of chlorous acid monitored in experiments over a range of reactant conditions was on the order of 10-4, as shown in Table 1.
expt no.a
[H2SO4] (M)
[NaClO3] (M)
[CH3OH] (M)
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17
4.50 4.50 4.50 4.50 4.50 4.50 4.50 4.50 4.50 4.50 3.25 3.25 3.25 3.25 3.25 3.25 3.25
0.05 0.10 0.20 0.30 0.40 0.40 0.20 0.20 0.20 0.20 0.80 0.80 0.80 0.40 0.75 0.75 0.75
0.23 0.23 0.23 0.23 0.23 0.23 0.05 0.10 0.30 0.40 0.23 0.23 0.23 0.23 0.15 0.15 0.15
a
titration (M × 104)
2.24
2.50 1.32 1.17
absorbance (M × 104) 0.70 1.05 2.01 2.46 3.35 2.70 1.00 1.00 2.00 2.26 2.52 2.54 1.50 1.40 2.00 2.10 4.50
Note: 1.75 M sodium sulfate was added in expt nos. 15 and
17.
It should be noted that the identities of the two intermediates were inconclusive and should be viewed in terms of the accuracy of the method of decoupling the absorbances. Chlorine dioxide has strong absorbances at all three wavelengths compared to those of the other two species, and therefore the correction of these signals was subject to considerable error. It is possible that the chlorine signal was spurious and that the initial signal may have actually represented another yet unidentified species. For example, in several experiments the reaction solution was monitored as a function of time at 370 nm, a wavelength at which chlorine has extremely low absorbance. Even at this wavelength, a significant maximum in the absorbancetime profile was observed. Also, at the concentrations we observed, chlorine would not have exhibited a yellow color that appeared and then disappeared in the initial phase of the experiments. The color was more characteristic of chlorine dioxide. The uncertainty in the identity of the initial signal and the potential error involved in decoupling the absorbances also cast doubt on whether chlorous acid was present or at least on the accuracy of chlorous acid concentrations determined in these experiments and reported in Table 1. Therefore, other techniques were required in order to verify the identities of species other than chlorine dioxide in these reaction experiments. The potentiometric titration procedure of Aieta et al. (1984) was employed to analyze chlorous acid in the solutions from several of the above batch reactor experiments. Details of the method are described by Indu (1993). Since chlorous acid established a steady-state concentration in these experiments, as shown in Figure 2, and in order to reduce the complexity of the analysis, the solutions were sampled after the initial chlorine peak had disappeared. Table 1 shows that chlorous acid was detected in those solutions that were investigated by the titration method. The table also shows that there was reasonable agreement between chlorous acid concentrations determined by the spectroscopic and titrimetric methods. (2) Experiments Incorporating Mercuric Nitrate. Further experiments were conducted using mercuric nitrate addition as a tool to test for the presence of chlorous acid and explore the impact of removing chlorous acid from reacting solutions. In one
Ind. Eng. Chem. Res., Vol. 36, No. 1, 1997 15
Figure 3. Influence of initially adding mercuric nitrate on the induction period in a methanol-chlorate-sulfuric acid batch reaction experiment at 3 M H2SO4, 0.24 M CH3OH, 2.4 M NaClO3, and 25 °C. Numbers represent the concentration of mercuric nitrate initially added.
set of experiments, mercuric nitrate was added at various concentrations to initial reaction solutions. Figure 3 shows chlorine dioxide-time profiles determined in the batch recirculation apparatus for several of these experiments all at the same reactant concentrations. The figure shows that, with increasing mercuric ion concentration, the duration of the induction period increased. In a related study, a mixture of 0.6 M sodium chlorate, 0.12 M methanol, 0.04 M mercuric nitrate, and 5.7 M sulfuric acid was prepared, reacted for 10 min, and then analyzed for oxidation products of methanol by proton NMR. The solution was found to contain formic acid. However, no chlorine dioxide formed. To determine whether mercuric ions oxidize methanol, another experiment was conducted at similar conditions but in the absence of sodium chlorate. The solution was analyzed by proton NMR after 10 min and found not to contain any traces of formic acid. (No formaldehyde was found in these solutions; however, we have confirmed in separate experiments that the reaction of formaldehyde with chlorate to formic acid is extremely fast under these conditions.) These NMR results showed that mercuric nitrate did not oxidize methanol to formic acid but that mercuric nitrate did not inhibit chlorate from oxidizing methanol. We suspect that formaldehyde is an intermediate in the process and that the steps to formic acid consist of eq 7 followed by,
H+ + HCHO + ClO3- f HClO2 + HCOOH
(18)
to yield the overall process
2H+ + CH3OH + 2ClO3- f 2HClO2 + HCOOH + H2O (19) There are two possible explanations for the influence of mercuric ions on the delay or absence of chlorine dioxide in these two sets of experiments: (1) chloride was an intermediate in the mechanism that formed chlorine dioxide; at the initial stages of the reaction, chloride ions formed by the sequential reduction of chlorate as described by eqs 1-3; mercuric ions sequestered these chloride ions, increasing the induction period, until all of the free mercuric ions had reacted; or (2) chlorous acid was an intermediate in the mechanism that formed chlorine dioxide; mercuric ions
Figure 4. Influence of adding mercuric nitrate on chlorine dioxide concentration in a methanol-chlorate-sulfuric acid batch reaction experiment (in progress) at 3 M H2SO4, 0.24 CH3OH, 2.4 M NaClO3, and 25 °C.
catalyzed the disproportionation of chlorous acid, inhibiting the formation of chlorine dioxide. Any chloride that was formed by chlorous acid disproportionation was sequestered by mercuric ions. As long as free mercuric ions were present, no chlorine dioxide formed; however, when all of the mercuric ions had reacted to form mercuric chloride, chlorine dioxide formation resumed. To explore these alternatives, we conducted two experiments in which mercuric nitrate was added to solutions in which the reactions had been initiated. In one experiment, a mixture of 0.12 M methanol and 0.6 M sodium chlorate in 5.7 M sulfuric acid was reacted in the batch reactor apparatus at 25 °C. Within 2 min, the UV absorbance at 370 nm increased from zero to 3 absorbance units and the solution became colored (yellow), indicating chlorine dioxide. At 2 min, mercuric nitrate was quickly added at a concentration of 0.04 M. The solution instantly became colorless, indicating that mercuric nitrate catalyzed the decomposition of chlorine dioxide. In a similar experiment, chlorine dioxide was allowed to form in a reaction mixture of 3 M sulfuric acid, 0.24 M methanol, and 2.4 M chlorate at 25 °C for 4 min, at which point 0.02 M mercuric nitrate was added. The chlorine dioxide concentration declined to zero as shown in Figure 4. After about an additional 7 min, chlorine dioxide started to form again. These latter results suggest that chlorous acid was present in the reaction mixtures prior to the addition of mercuric ions. We have shown earlier in this study that mercuric ions enhanced the rate of disproportionation of chlorous acid, enhanced the rate of chlorine dioxide disappearance when both chlorous acid and chlorine dioxide were present, but did not decompose chlorine dioxide in the absence of chlorous acid. Additionally, the behavior shown in Figure 4 cannot be explained if the only effect of mercuric nitrate addition is sequestering of chloride ions, since the removal of chloride ions would not have caused the chlorine dioxide, already formed, to decompose. We briefly explored the use of formic acid in the reduction of chlorate ions. We found that the rate of this reaction was much slower than that of the corresponding methanol-chlorate reaction. We also found that mercuric ions have almost no effect on the rate of the formic acid-chlorate reaction, suggesting that chlorous acid is not an intermediate in that reaction.
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Ind. Eng. Chem. Res., Vol. 36, No. 1, 1997
Conclusions Mercuric ions appear to catalyze the disproportionation of chlorous acid in aqueous solution but do not decompose chlorine dioxide in solution unless chlorous acid is also present. In solutions containing both chlorous acid and chlorine dioxide, mercuric ions appear to catalyze the decomposition of both species. This behavior has provided a qualitative means of demonstrating the presence of chlorous acid during the formation of chlorine dioxide by the reaction of methanol and chlorate ions. Acknowledgment We gratefully acknowledge financial support from Akzo Nobel for this work. We thank Michael Burke, Helena Falge´n, John Gray, and Joel Tenney of Akzo Nobel for helpful discussions regarding industrial chlorine dioxide processes. We are grateful to Professor Gilbert Gordon, Chemistry Department, Miami University, Oxford, OH, for reading the manuscript and for his suggestions regarding the possible role of mercuric ions in the decomposition of chlorous acid, which we have incorporated in the paper. Literature Cited Adam, L. C.; Fabian, I.; Suzuki, K.; Gordon, G. Hypochlorous Acid Decomposition in the pH 5-8 Region. Inorg. Chem. 1992, 31, 3534. Aieta, E. M.; Roberts, P. V.; Hernandez, M. Determination of Chlorine Dioxide, Chlorine, Chlorite, and Chlorate in Water. J. Am. Water Works Assoc. 1984, Jan, 64. Burke, M.; Tenney, J.; Indu, B.; Hoq, M. F.; Carr, S.; Ernst, W. R. Kinetics of Hydrogen Peroxide-Chlorate Reaction in the Formation of Chlorine Dioxide. Ind. Eng. Chem. Res. 1993, 32, 1449. Ciavatta, L.; Grimaldi, M. The Hydrolysis of Mercury(II) Chloride, HgCl2. J. Inorg. Nucl. Chem. 1968, 30, 563. Coetzee, J. F.; Ritchie, C. D. Solute-Solvent Interactions; Marcell Dekker: New York, 1969; p 372.
Ernst, W. R.; Indu, B.; Crump, B.; Gelbaum, L. T. Kinetics of the Reaction of Methanol with Chlorate Ions in Acidic Solution by Proton-NMR Spectrometry. AIChE J. 1996, 42, 1379. Fa´bia´n, I.; Gordon, G. Iron(III)-Catalyzed Decomposition of the Chlorite ion: An Inorganic Application of the Quenched StoppedFlow Method. Inorg. Chem. 1992, 31, 2144. Gordon, G.; Emmenegger, F. Complex Ion Formation Between Chlorine Dioxide and Chlorite. Inorg. Nucl. Chem. Lett. 1966, 2, 395. Hong, C. C.; Lenzi, F.; Rapson, W. H. The Kinetics and Mechanism of the Chloride-Chlorate Reaction. Can. J. Chem. Eng. 1967, 45, 349. Hoq, M. F.; Indu, B.; Ernst, W. R.; Neumann, H. M. Kinetics of the Reaction of Chlorine with Formic Acid in Aqueous Sulfuric Acid. J. Phys. Chem. 1991, 95, 681. Indu, B. Kinetics of Methanol-Chlorate Reaction in the Formation of Chlorine Dioxide. Ph.D. Dissertation, Georgia Institute of Technology, Atlanta, GA, Aug 1993. Indu, B.; Hoq, M. F.; Ernst, W. R. Acidity of Sulfuric Acid-Sulfate Solutions By Kinetic Measurements. AIChE J. 1991, 37, 1744. Kieffer, R. G.; Gordon, G. Disproportionation of Chlorous Acid: Part 1. Stoichiometry. Inorg. Chem. 1968, 7 (2), 235. Lenzi, F.; Rapson, W. H. Further Studies of the Mechanism of Formation of Chlorine Dioxide. Pulp Pap. Mag. Can. 1962, Sept, T-442. Masschelein, W. J. Chlorine Dioxide; Ann Arbor Science: Ann Arbor, MI, 1979. Schmitz, G.; Rooze, H. Mechanism for Reactions of Chlorite and Chlorine Dioxide. 2. Kinetics for Reactions of Chlorite in the Presence of o-Toluidene. Can. J. Chem. 1984, 62, 2231. Tenney, J.; Shoaei, M.; Obijeski, T.; Ernst, W. R.; Lindstroem, R.; Sunblad, B.; Wanngard, J. An Experimental Investigation of the Chloride-Chlorate Reaction System. Ind. Eng. Chem. Res. 1990, 29, 916.
Received for review July 22, 1996 Revised manuscript received October 21, 1996 Accepted October 21, 1996X IE960431J
X Abstract published in Advance ACS Abstracts, December 1, 1996.
