Environ. Sci. Technol. 1984, 18, 212-216
Edginpton, D. N.; Alberta, J. J.; Walhgren,M. A.; Karttunen, J. 0.; Reeve, C. A. “Transuranium Nuclides in the Environment, Proceedings of a Symposium”; International Atomic Energy Agency: Vienna, Austria, 1976; pp 493-516. Nelson, J. L.; Perkins, R. W.; Nielsen, J. M.; Haushild, W. L. “Disposal of Radioactive Wastes into Seas, Oceans, and SurfaceWaters, Proceedings of a Sympmium”;International Atomic Energy Agency: Vienna, Austria, 1966; pp 139-161. Perkins, R. W.; Nelson, J. L.; Haushild, W. L. Limnol. Oceanogr. 1966, 11, 235-248.
Beasley, T. M., Oregon State University, Corvallis, OR, unpublished data, 1983. Gillham, R. W.; Cherry, J. A.; Lindsay, L. E. Health Phys.
(27) U.S. Energy Research and Development Administration, New York, 1977, Health and Safety Laboratory Quarterly Report HASL-329, pp 1-401. (28) Drever, J. I. “The Geochemistry of Natural Waters”; Prentice-Hall: Englewood Cliffs, NJ, 1982; Chapter 1. (29) Sprugel, D. G.; Bartelt, G. E. J . Environ. Qual. 1978, 7 (2), 175-177. (30) Hayes, D. W.; Horton, J. H. In “TransuranicElements in
the Environment”;Hanson, W. C., Ed.; Technical Information Center: Springfield, VA, 1980; pp 602-611. (31) National Bureau of Standards, Standard Reference Material 4350B, River Sediment, 1982.
1980,39,637-649.
Knebel, H. J.; Kelley, J. C.; Whetten, J. T. J . Sediment. Petrol. 1968, 38, 600-611.
Whetten, J. T.; Kelley, J. C.; Hanson, L. G. J . Sediment. Petrol. 1969, 39, 1149-1166. Baker, E. T. Geol. SOC.Am. Bull. 1976,87, 625-632.
Received for review February 18, 1983. Revised manuscript received August 11, 1983. Accepted September 12, 1983. Our research is funded by the Environmental Research Division, Office of Health and Environmental Research, US.Department of Energy.
Stoichiometry of the Reaction between Chlorite Ion and Hypochlorous Acid at pH 5 Tsung-fel Tang and Gilbert Gordon” Department of Chemistry, Miami University, Oxford, Ohio 45056
The availability of an improved analytical method allowed a preliminary study of the stoichiometry of the reaction between chlorite ion and hypochlorous acid at pH 5. The stoichiometry varies with the initial concentration ratio of the reactants, suggesting parallel pathways for the reaction mechanisms. The yield of chlorine dioxide is shown to vary with the initial concentration ratio of hypochlorous acid to chlorite ion as well as the contact time between reactants. The formation of molecular oxygen is postulated as a byproduct.
Introduction The major method used by water treatment plants for the generation of chlorine dioxide is mixing aqueous solutions of chlorine and sodium chlorite (1-4).Granstrom and Lee (2) surveyed 56 water treatment plants using chlorine dioxide. The ratio of chlorine to sodium chlorite used by these water treatment plants for the generation of chlorine dioxide varied from 1 0 1 to 1:l. The survey also indicated only 25 plants measured the pH of the solution and that the effluent pH of the chlorine dioxide generator varied from 1.2 to 10.4. The present experimental results demonstrate that chlorate ion is generated at pH 5 and that side reactions also consume chlorine dioxide and form oxygen as well as chlorate ion. The stoichiometric details of this reaction above pH 5 are not clearly understood since other products in addition to chlorine dioxide may be formed (2). In order to study the overall reaction, it is necessary to determine quantitatively the various oxychlorine species, namely, hypochlorous acid (hypochlorite ion), chlorite ion, chlorine dioxide, and chlorate ion. The analytical determination of these oxychlorine species is complicated due to their similar chemical and physical properties. Chlorine hydrolyzes in aqueous solution at pH 5 to form hypochlorous acid. A preliminary kinetic study of the reaction between chlorite ion and hypochlorous acid was carried out by Granstrom et al. (5). However, they re212
Envlron. Sci. Technol., Vol. 18, No. 3, 1984
ported many experimental difficulties were involved. The analytical method they used for the determination of chlorine dioxide, chlorite ion, and hypochlorous acid was the spectrophotometric method, but due to the overlap of the various spectra in the UV region, the computation for the concentrations of the oxychlorine species necessitated the solution of three simultaneous equations. A slight error in the value of the molar absorptivity or of the observed absorbance is magnified considerably in the final values obtained for the concentrations of hypochlorous acid, chlorite ion, and chlorine dioxide. Another problem associated with the analytical method reported by Granstrom et al. (5) is the determination of total chlorine in the solution. The total chlorine was determined by passage of the sample through a Jones reductor, and the resulting chloride ion formed was titrated potentiometrically with silver nitrate. However, the high volatility of chlorine dioxide increased the difficulty of sample handling, and some chlorine dioxide invariably was lost to the atmosphere (5). Thus, an error of nearly 5% was reported for the total chlorine determination. In the present study, the stoichiometry of the reaction between chlorite ion and hypochlorous acid a t pH 5 was determined under various reactant ratios. The experimental difficulties described by Granstrom were minimized by using an improved analytical procedure (6). The procedure involves the combination of two analytical methods, namely, the ion selective electrode method (6) and the iodometric method (1, 7). The iodometric method is used to determine the oxychlorine species in terms of redox properties, and the ion selective electrode method is used to determine the chloride ion, chlorate ion, and the total chlorine by the halogen content. Thus, the hypochlorous acid-chlorite ion reaction can be studied more precisely by utilizing two different properties.
Experimental Section
Reagents. Reagent-grade chemicals and triply deionized distilled water were used throughout. Preparation and
0013-936X/84/0918-0212$01.50/0
0 1984 American Chemical Society
Table I. Replicate Sample Analysis of the Reaction with Chlorite Ion Present in Excess reactants
products"
M X 103 N X 103
c1HOC1 ClO;
2.26 f 0.02 2.37 f 0.01 6.62 f 0.02
NX
103
5.15 f 0.03
4.74 f 0.02 26.48 f 0.08
c10, (30,-
total
M X 103
11.25 f 0.04
31.22 f 0.09
I/
i
1.61's O.O1,b 6.44 f 1.60 f 0.01' 0.04 4.40 ?: 0.01 22.00 f 0.05 0.06 's 0.03 0.36 f 0.18 11.22 f 0.05 28.8 's 0.2
3.01
(1
E 2 X
I 0-
The species involved in the products were analyzed 1 6 Result E obtained by h after initiation of the reaction. the iodometric method. ' Result C/4 obtained by the ion selective electrode method. a
1
7.01
5.0
0
1
c1-
5
10
15 20 time, hours
25
30
Figure 2. Reactant and product concentrations as a function of time ([hypochlorite],,,,,/[chiorite],,,,, = 1.30).
Table 11. Product Analysis of the ClO,--HOCI Reaction at Various Times with Excess Chlorite Ion" 2.0
time, min
M .. -.. -.. r..
0
5
10
15
20
CIOj -..-..-..-.,-.I
25
30
time, hours
Flgure 1. Reactant and product concentrations as a function of time ([chlorRe],,,,/[hypochlorite],,,, = 3.46).
standardization of the sodium thiosulfate solution, sodium chlorite solution, osmium tetraoxide solution, and silver nitrate solution in 70% methanol were described elsewhere (6,8).Acetate buffer a t pH 5.03 (0.1 M) was prepared by dissolution of the appropriate amounts of sodium acetate and acetic acid in triply distilled water. Hypochlorous acid was prepared by the method of Cady (9). Chlorine monoxide is produced by the reaction of the yellow form of mercuric oxide with chlorine dissolved in carbon tetrachloride. The hypochlorous acid solutions were obtained by shaking an ice-cold 0.05 N HCIOl solution with the chlorine monoxide solution. Hypochlorous acid solutions were prepared fresh daily and standardized iodometrically by using a shrinking bottle (10) as the storage container. An aliquot of hypochlorous acid solution was transferred directly from the shrinking bottle to a Silverman syringe (11)via a glass adaptor which was placed under the surface of a potassium iodide solution. This transfer procedure avoided the loss of chlorine and markedly increased the precision of the procedure. Instrumentation. The potentiometric titration system included a Metrohm Herisau Model E536 potentiograph, a Metrohm Herisau Model E535 autoburet, and an Orion Model 96-17 combination ion selective electrode. A Metrohm Herisau E535 autoburet was used to deliver the sodium thiosulfate titrant to the Thyodene (a stable, soluble Fisher dry starch-type indicator) end point (12). Procedures. The reaction vessel for the stoichiometry
70 500 1920
[HOClI, [ClO,-I, [Cl-I, [ClO,], mM mM mM mM 0 0 0
2.71 2.59 2.33
3.98 4.15 4.45
3.76 3.70 3.54
[C103-l, mM 0 0 0.18
" Conditions: [ C l - ] , i ~ = 1.78 x M, [HOClIjrits = 1.93 X M, and [ C l O , - ] ~ =i ~6.68 X M in 0.1 M acetate buffer a t pH 5.03. study was a 100-mL syringe. The volatility of chlorine dioxide made it necessary to use a syringe vessel (i.e., a shrinking bottle (10)) from which any air gap could be removed; otherwise, the chlorine dioxide concentration in the solution would be dependent on the size of the air gap, due to the volatile chlorine dioxide being equilibrated with the air. At the beginning of each experiment, the reaction syringe was clamped in an inverted position, and exact volumes of the acetate buffer, chlorite ion, chloride ion, and hypochlorous acid were delivered to the syringe by utilizing Silverman syringes (11). Any air gap was quickly removed, and the syringe was inverted and tightly sealed with a Teflon cap during the course of the reaction. No chlorine dioxide loss was observed during the reaction. The sample was transferred directly from the reaction syringe to a Silverman syringe via a glass adaptor; thus, no chlorine dioxide was lost during the transferring process, and no air gap was formed in the reaction syringe after the sample had been transferred. In a typical experiment, an aliquot of the sample was quenched with potassium iodide solution a t pH 5 and acidified to pH 2.3 by adding 1 M hydrochloric acid. The iodine formed was titrated with standard sodium thiosulfate solution. Thus, the s u m total normality of chlorine dioxide, chlorite ion, and hypochlorous acid can be calculated from this titration (result A). Another aliquot was adjusted to pH 7.8 by adding pH 8.2 dihydrogen phosphate buffer. The resulting solution was purged with nitrogen gas for 5 min to eliminate Environ. Sci. Technoi., Vol. 18, No. 3, 1984
213
Table 111. Stoichiometry of the CIO,--HOCl Reactiona, with Chlorite Ion Present in Excess
time, min
n/m
Plm
q/m
rlm
stexcess n
s/r
oxidizing power recovered,c %
70 2.06 1.13 0 1.95 0.09 0.06 96.9 500 2.12 1.23 0 1.92 0.17 0.12 94.4 1920 2.25 1.38 0.09 1.83 0.24 0.16 92.7 a The reaction is represented by the equation mHOCl + nC10; = pCl- + qC10,- + rC10, + SO,. Reaction conditions: [C~-]WW= 1.78 X lo'' M, [ H O C l l h l ~= 1.93 X l o w 3M, and [ClO,']wu = 6.68 X lo-' M in 0.1 M acetate buffer at pH 5.03. Ratio of (total product oxidizing power recovered/total reactant oxidizing power) X 100. Table IV. Product Analysis of the CIO,'-HOCl Reaction at Various Times with Excess Hypochlorous Acida time, [HOCll, [ClO,'], min mM mM
[Cl-I, mM
[ClO,], mM
[ClO;], mM
130 420 1540
4.15 4.41 4.90
1.88 1.53 0.69
0.30 0.53 1.17
a Conditions: [ C l - ] h i ~ = 2.75 X lo-' M, [HOClIipim = 3.01 X M, and [ClO,-]fi, lo-' M in 0.1 M acetate buffer at pH 5.03.
= 2.31
1.65 1.60 1.17
0.09 0.09 0.09
Results and Discussion By use of the results of the iodometric method, the molar concentrations of hypochlorous acid, chlorite ion, and chlorine dioxide are calculated by using the following equations: [HOC11 = B/2
(1)
[ClOz-1 = c / 4
(2)
Environ. Sci. Technol., Vol. 18, No. 3, 1984
[Cl-] = D
(4)
[ClO2-] = E
(5)
[ClO3-] = F - ( A / 5 ) - (3B/10) - (C/20) - D
(6)
x
chlorine dioxide. Then, several crystals of potassium iodide were dissolved in this chlorine dioxide free solution, and only hypochlorous acid reacted with iodide ion a t pH 7.8. The iodine formed was titrated with standard sodium thiosulfate solution, and the result corresponded to the normality of hypochlorous acid (result B). The solution was further acidified to pH 2.3 by adding 1 M hydrochloric acid. The iodine formed was again titrated with sodium thiosulfate solution, and the result corresponded directly to the normality of chlorite ion (result C). In the case of the reaction between excess chlorite ion and hypochlorous acid, the concentrations of chloride ion (result D ) and chlorite ion (result E ) were determined by the ion selective electrode method. Another aliquot of sample was quenched (5, 7) with As(III), and 3 drops of 0.1% OsOl in 1 M sulfuric acid was added. The resulting chloride ion formed was titrated with silver nitrate in 70% methanol, and the result corresponded to the total chlorine (result F). In the case of the reaction of excess hypochlorous acid with chlorite ion, only chloride ion and total chlorine concentration were determined by the ion selective electrode method, since the hypochlorous acid was found to interfere with the chlorite ion determination (8). In a series of preliminary experiments in the presence of excess chlorite ion ([C102-]/[HC10] > 2), replicate samples were analyzed by the above procedure, and the precision of the analytical method was established to be better than 0.5%. The concentrations of the reactants and the products were also determined as a function of time in the presence of either excess chlorite or hypochlorous acid with the same precision and accuracy.
214
where result A is the total normality of chlorine dioxide, hypochlorous acid, and chlorite ion, result B is the normality of hypochlorous acid, and result C is the normality of chlorite ion. By the ion selective electrode method, the concentration of chloride ion, chlorite ion, and chlorate ion can be calculated by the following equations:
where result D is the chloride ion concentration, result E is the chlorite ion concentration, and result F is the total chlorine contained in the sample. Since the chlorite ion concentration can be obtained by both methods, results E and C/4 should be in agreement within experimental error. The stoichiometry was studied by analyzing replicate samples in the presence of excess chlorite ion ([C102-]/ [HOC11 = 2.79). The results are shown in Table I, where the uncertainties of the results represent the standard deviation from the mean of four replicate samples. It is noted that the results for the chlorite ion determination are in good agreement between the two methods; i.e., results E and C/4 are in agreement. The chlorine balance between reactants and products can be calculated by using the data given in Table I. The total chlorine available in M, and the total the reactants is (11.25 f 0.04) X chlorine found in the products is (11.22 f 0.05) X M. This represents a chlorine recovery of 99.3%which implies that there is no volatile chlorine species loss either from the reaction syringe or from the sample transferring process and that chlorine species other than those presented in Table I are not formed. On the other hand, the redox balance can be calculated by adding the total normality of the reactants, (31.22 f 0.09) X N, and total normality of the products, (28.8 f 0.2) X N. Clearly, only 93% of the total oxidizing power is found in products, which strongly suggests that after 16 h-in addition to oxychlorine species-there is some other oxygen-containing product formed. It was also observed that gas bubbles formed in the reaction syringe even though it was sealed very tightly during the reaction. The gas bubbles are indicative of the formation of oxygen, and it would explain the oxidizing power deficiency (8). The stoichiometry was studied further by determining the concentration of each species involved in the reaction as a function of time. Figure 1 demonstrates that chlorine dioxide forms rapidly and that all of the hypochlorous acid is consumed in the presence of excess chlorite ion ([ClO,]/ [HOC11 = 3.46). Chlorate ion was not detectable a t the beginning of the reaction, but its concentration increased slowly. This suggests that in this case the chlorate ion formation is due to slow reactions which involve chlorine
*
Table V. Stoichiometry of the (30,'-HOCI Reactiona? with Hypochlorous Acid Present in Excess
time, min
n/m
P/m
4/m
r/m
130 420 1540
1.63 1.57 1.20
1.03 1.18 1.1 8
0.22 0.37 0.64
1.38 1.09 0.38
slexcess m
s/rd
oxidizing power recovered: %
0.06 0.14 0.44
0.05 0.14 0.74
97.0 92.6 84.0
a The reaction is represented by the equation mHOCl + nC10,' = p a - + qC10; t rC10, t SO,. Reaction conditions: M in 0.1 M acetate buffer at [Cl']h%g= 2.75 X M,[HOCl]hi++ = 3.01 x M, and [ C l O , ' ] ~ u= 2.31 X Determined from pH 5.03. Ratio of (total product oxidizing power recovered/total reactant oxidizing power) X 100. oxidizing power deficiency.
dioxide and chlorite ion. According to Emerich and Gordon (13,14),the decomposition of chlorine dioxide forms chlorate ion and chlorite ion: 2C102 + H 2 0
-
C103- + C102- + 2H+
(7)
also, Buser and Hanisch (15)report the decomposition of neutral solutions of chlorite ion produces chloride ion and chlorate ion: 3C102-
-
2C103-
+ C1-
(8)
Both reactions 7 and 8 are reported to be slow a t pH 5, which is in agreement with the present observations. Figure 2 demonstrates that in the presence of excess hypochlorous acid, there is also a rapid reaction forming chlorine dioxide, and nearly all of the chlorite ion is consumed rapidly. A comparison of Figures 1 and 2 shows that there is more chlorate ion formed a t the beginning of the reaction, and the chlorate ion increases considerably slower in Figure 1 than in Figure 2, where hypochlorous acid is in excess. The chlorate ion formation in Figure 2 can be interpreted in terms of the reaction between chlorine dioxide and excess hypochlorous acid (16,17) HOCl
+ 2C102 + H 2 0
-
2C103-
+ C1- + 3H+
(9)
and the decomposition of hypochlorous acid (18) which is in excess as is noted below: 3HOC1- 3H+
+ 2C1- + C103-
(10)
Burghardt (17)also postulates the formation of chlorine (C12) which could result from the rapid, direct reaction between HOCl and the C1- produced. At this point, the reaction between chlorite ion and hypochlorous acid (and/or chlorine) a t pH 5 can be summarized in terms of the equation mHOCl
+ nC102-
-
pC1-
+ qC103- + rCIOB (11)
The coefficients m, n, p, q, and r can be obtained directly from the experimental results. The data in Tables I1 and I11 demonstrate that in the presence of excess chlorite ion eq 11 can be approximated by the equation HOCl
+ 2C109,- + H+
-
C1-
+ 2C102 + H2O
(12) However, the stoichiometry reported in Table IV demonstrates that in the presence of excess hypochlorous acid, chlorate ion is present as one of the initial products. Thus, the overall reaction deviates from the simple process shown by eq 12. It is also noted from the results shown in both Tables I11 and V that the percentage of total oxidizing power recovered in the products decreases but the ratio between oxygen and excess reactant or chlorine dioxide increases with respect to the reaction time. These results clearly suggest that the formation of oxygen is due to side reactions (19,20).Flis reports (19)that reaction 13 gen2c102 Cl2 202 (13)
-
+
erating oxygen is operative a t pH 5. In the presence of excess hypochlorous acid, the hypochlorous acid (and/or chlorine) could also decompose (18)slowly as described by the equation 2HOC1- 2HC1+ O2
(14)
However, in the presence of excess chlorite ion a t pH 5, the major decomposition product of chlorite ion is chlorine dioxide, but approximately 3% of the chlorite ion is converted to oxygen (20)as illustrated by the equation ClO2-
- c1- +
02
(15)
Emmenegger and Gordon studied the reaction between sodium chlorite and dissolved chlorine in 0.2 M perchloric acid (21)and noted that with chlorine(II1)in excess, there was marked improvement in the production of chlorine dioxide. They also noted that in more dilute solutions with a constant chlorine(0) to chlorine(II1)ratio (1:2 mole ratio), the production of chlorate ion is favored. In the present study a t pH 5, excess chlorite ion also gives marked improvement in the production of chlorine dioxide as is summarized in Tables I11 and V. The changing stoichiometry reveals the existence of parallel pathways that contribute differently depending on the initial concentration ratios (22).Therefore, additional experiments such as pH-dependent studies and kinetic studies are necessary to demonstrate the microscopic details of these parallel reaction pathways. Conclusion The present study demonstrated the use of newly developed analytical procedures to study the reaction between hypochlorous acid and chlorite ion at pH 5 in an acetate buffer. The experimental results reveal that parallel pathways are involved in the reaction since the stoichiometry varies with the initial concentration ratio of the reactants. Specifically, in the presence of excess hypochlorous acid, reduced chlorine dioxide yields are observed. Clearly, the experimental results also indicate the importance of the initial concentration ratio of hypochlorous acid (and/or chlorine) to chlorite ion as well as the contact time between reactants for water treatment plant using chlorine dioxide. Registry No. C102-, 14998-27-7; HOC1, 7790-92-3.
Literature Cited (1) White, G. C. "Handbook of Chlorination";Van Nostrand Reinhold: New York, 1972. (2) Granstrom, M. L.;Lee, G. F. J.-Am. Water Works Assoc. 1958,50, 1453-1466. (3) Aieta, E. M.; Roberta, P. V. J.-Am. Water Works Assoc. 1984, 76, 64. (4) Denis, M.; Masschelein, W. J. Analusis 1983, 11, 79. Envlron. Sci. Technol., Vol. 18, No. 3, 1984
215
(5) Granstrom, M. L.; Israel, B. M.; Snow, W.; Lynch, M. Rutgers, The State University of New Jersey, Sanitary Engineering Laboratories, unpublished results, 1964. (6) Tang, T.-F.; Gordon, G. Anal. Chem. 1980,52,1430. (7) Hong, C. C.;Rapson, W. H. Can. J. Chem. 1968,46,2061. ( 8 ) Tang, T.-F. M.S. Thesis, Miami University, Oxford, OH, 1980. (9) Cady, G. H. Inorg. Synth. 1957,5, 156. (10) Silverman, R. A.; Gordon, G. Anal. Chem. 1974,46,178. (11) Silverman, R.A.; Gordon, G. J. Chem. Educ. 1973,50,654. (12) Walker, G. T.; Freeman, F. M. Seifen, Ole, Fette, Wachse 1956,82,187. (13) Emerich, D.; Gordon, G. “Abstracts of Papers”, 179th National Meeting of the American Chemical Society, Houston, TX, March 23-28, 1980;American Chemical Society: Washington, DC, 1980. (14) Emerich, D. Ph.D. Dissertation, Miami University, Oxford, OH, 1981.
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Environ. Sci. Technol., Vol. 18, No. 3, 1984
(15) Buser, W.; Hanisch, H. Helv. Chim. Acta 1952,35,2574. (16) White, J. F.;Taylor, M. C.; Vincent, G. P. Znd. Eng. Chem. 1942,34,782. (17) Burghard, S. I. Northwest Sci. 1951,25, 137. (18) Wojbwicz, J. A. In “Kirk-Other Encyclopedia of Chemical Technology”, 3rd ed.; Interscience, New York, 1980. (19) Flis, I. E.; Mischenko, K. P.; Salnis, K. Y. Zh.Prikl. Khim. (Leningrad) 1959,32,284. (20) Gordon, G.; Kieffer, R. G.; Rosenblatt, D. H. Prog. Znorg. Chem. 1972,15,208. (21) Emmenegger, F.; Gordon, G. Znorg. Chem. 1967,6,633. (22) Gordon, G. “Understanding Chemical Kinetics and Interpreting Mechanisms in Solution”;Department of Chemistry, Miami University: Oxford, OH, 1977.
Received for review July 22,1983. Accepted October 11, 1983.
