Environ. Sci. Technol. 1988, 22, 1171-1 177
Fate of Sulfur( I V ) Dechlorinating Agents in Natural Waters: Effect of Suspended Sedimentst Phltlp J. Kijak" and George R. Helz
Department of Chemistry and Biochemistry, University of Maryland, College Park, Maryland 20742 To investigate the fate of SOz in dechlorinated waste water effluents, oxidation rates were measured in nonilluminated solutions at near-neutral pH and 25 OC. River water was simulated with 0.01 M NaC1, 0.001 M buffer, and 1 g/L standard sediment MESS-1. Components leached from the sediment catalyzed the oxidation of S(IV) (i,e,, HS03- S03z-.) by 02,but the particles themselves exerted a slight inhibitory effect. Sulfate was the major reaction product. Some nonoxidative loss of S(1V) to particles was observed at high-sediment concentrations (20 g/L). Sulfur(IV)reductively dissolved 25% of the Cu from the sediment, possibly an environmentallyharmful process. Iron and manganese dissolutions were insignificant. The rate of loss of S(1V) from air-saturated solutions covering a 50-fold S(1V) concentration range was well described by the empirical equation (time in s and concentrations in M) -d[SO,'-]/dt = (5 X lO-')[[S(IV)]/(l + [Ht]/Ka)]1/2 K, being the second ionization constant of HzSO,. The rate of loss of S(1V) was a factor of 2 faster in actual effluent/river water mixtures, likely caused by higher trace metal concentrations in these mixtures.
Table I. Characteristics of MESS-1 elemental composition"
c, %
+
Introduction Because chlorinated waste water effluents may pose a threat to aquatic biota (1, 21, treatment plant operators in some parts of the United States are installing dechlorination systems. Dechlorination is usually accomplished by injection of excess sulfur(1V) compounds (Le., SO2, bisulfate salts, or sulfite salts) after the waste water has been disinfected with chlorine. Virtually nothing is known about the fate of the S(1V) compounds in receiving waters. Sulfur dioxide and its hydrolysis products are not innocuous to organisms and biological compounds. Sulfur dioxide has been used for decades to control food and wine spoilage. Metal-catalyzed S(1V) oxidation has been shown to bring about destruction of amino acids (3) and DNA (4). Oxidation of S(IV)at high pH produces SOs%in significant yields (5). Mixtures of SOsz-and bromide, which is an important component of sea salt, have been patented as a disinfectant (6). Because S(1V) compounds, especially in the presence of oxygen, may have biological effects, it is desirable to know something about their lifetime and chemical fate in the aquatic environment. Sulfur(1V) dechlorinating agents are expected to disappear from natural waters by oxidation processes leading to thermodynamically stable sulfate. On the basis of geochemical abundances (7),dissolved oxygen or particulate iron and manganese oxyhydroxides are probably the only oxidizing agents of potential significance. Other oxidative species, such as Fe3+,CuZt, HAs04*, Cr04z-,IO3-, etc., will normally be present in rivers at concentrations that are negligible relative to S(1V) concentrations, which typically are 10-5-10-4 M in dechlorinated effluents. *Address correspondence to this author at the Food and Drug Administration, Division of Veterinary Medical Research, Building 328-A, BARC-East, Beltaville, MD 20706. +This paper is from a dissertation submitted to the Graduate School, University of Maryland, by P.J.K. in partial fulfillment of the requirement for the Ph.D. degree in Chemistry. 0013-936X/88/0922-1171$01.50/0
2.99 f 0.09
11.03 f 0.38 A1 as A1,0,, % 67.5 f 1.9 Si as SiO,, % 4.36 f 0.25 Fe as FezOs, % $13 f 25 Mn, I.Lg/g 10.8 f 1.9 COT rg/g 25.1 f 3.8 cu, P P l P 34.0 f 6.1 Pb, I.Lg/g major mineral phases by X-ray diffractionb a-quartz
Elemental composition data taken from ref 11. Several peaks in the diffraction pattern indicated a feldspar-type of mineral is present, but not enough detail was present to positively identify a specific mineral phase.
Thermodynamically, particulate iron/manganese oxyhydroxides are capable of oxidizing S(1V) to sulfate at acidic to neutral pH values (8), and this process has in fact been observed under acidic conditions (9). The reaction leads to reductive dissolution of Fe and Mn. In rivers, this process could result in the release of coprecipitated trace metals, such as copper, in environmentally significant quantities. The purpose of the research reported here is to explore the role of suspended sediments containing iron/manganese oxyhydroxides as (a) a source of oxidizing capacity and (b) a catalyst or inhibitor in the Ozoxidation of S(1V). We have performed a series of experiments in which sulfite loss and sulfate gain were observed in suspensions containing 1g/L standard sediment, MESS-1, issued by the National Research Council of Canada. Highly turbid river waters, such as the Mississippi, commonly contain up to 1 g/L suspended sediment, although most river waters contain less, except at flood stage (10). MESS-1 is an estuarine sediment prepared simply by freeze-drying, screening, and sterilizing by radiation. No heat treatment or grinding, either of which could modify natural mineral phases, was used (11). Selected chemical characteristics of MESS-1 are shown in Table I. Light was excluded from our test solutions to prevent possible photochemical reactions between S(1V) and iron oxyhydroxides (12). The effect of pH was studied by doing experiments at pH 7.1 and pH 8.4, values that fall within the range of pH in river waters. As we will show, solutions equilibrated for 2 days with MESS-1 achieve solute compositions similar in several respects to average river water. Experimental Procedures Apparatus. The experimentswere done in four-necked, 1-L round-bottom flasks. The flasks were wrapped with opaque material to exclude light. Temperature was kept at 25 "C by placing the flasks in a temperature-controlled water bath. The water-sediment suspension was stirred constantly. All the syringes, bottles, glassware, and filter holders were washed with soap and water, rinsed, and then soaked in 5% nitric acid for a minimum of 24 h. They were rinsed twice with deionized water, soaked in deionized water for a minimum of 24 h, and then rinsed twice more. The
0 1988 American Chemical Society
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washing procedure removed all Fe, Mn, Co, Cu, and P b contamination detectable in blanks by graphite furnace atomic absorption spectrometry, and the thorough rinsing procedure removed all traces of nitrate, which interfered in the sulfite analysis. Solutions. Solutions were prepared from 10 MQ or better of deionized water (Millipore Milli-Q system) to which was added reagent-grade NaCl to control ionic M. Solutions were buffered with strength at 1 X M of either Mops [3-(morpholino)propanesulfonicacid] or Tris [tris(hydroxymethyl)aminomethane]. Mops controlled pH a t 7.1, whereas Tris controlled it at 8.4. Procedure. Prior to the addition of sulfite, the sediment suspension was stirred for 2 days to equilibrate the sediment with the solution. Two days were sufficient for the dissolved trace metal concentrations to stabilize. Then samples were taken to obtain the initial concentrations of the metals and sulfate. Air saturation was maintained during the course of these experiments by continuously bubbling air through the flask. In those runs where deaerated conditions were desired, N2 was used in place of air, and the flask was maintained in a N2-filledglovebox. The O2concentration was monitored with a Clark-type electrode. Samples were taken with a 50-mL plastic syringe connected to a plastic tube that could be inserted into the flask. The mixtures were filtered through 0.2-pm Nuclepore filters. A total of 4 mL was transferred through the filter into a second syringe, avoiding air contact. This subsample was immediately analyzed for sulfite and sulfate. The rest of the sample was filtered into a polypropylene bottle, acidified with Ultrex nitric acid (J.T. Baker Chemical Co.), and analyzed later for trace metals. Analysis. In the field and in experiments with waste water effluent, sulfite was measured by iodometric titration with amperometric end point detection. Blanks containing no sulfite suggested that interference from other iodinereactive substances was small and constant. In all the laboratory experiments with sediment suspensions and filtrates, sulfate and sulfite were determined by ion chromatography. A Dionex QIC IC equipped with two Dionex HPIC-AG1 guard columns in series and a Dionex ASF fiber suppressor were used for the analysis. A 0.003 M NaHC03/0.0024 M Na2C03eluent was used. The fiber suppressor regenerant was 0.024 M H2S04 as recommended by the manufacturer. The flow rate was 2.5 mL/min. Under these conditions the sulfite peak eluted in 6 min and the sulfate peak in 10 min. The two peaks were completely separated. No evidence was found that the sulfite was oxidizing to sulfate within the instrument, as has been reported (13). Duplicate determinations of sulfite by ion chromatography agree within 5 % , although larger discrepancies occurred in about 12% of the cases. Duplicate determinations of sulfate generally agreed within 10%. In duplicate sulfate determinations, a tendency was noted for the second measurement to exceed the first by about 5%. Most of the sulfite and sulfate concentrations in this paper represent means of duplicate determinations, the main exception being the data collected in the first 15 min of each run. Sulfite standards were freshly prepared as needed from reagent-grade Na2SOs. In the early stages of the work, the standards were calibrated by iodometric titration with amperometric end point detection. However, this was found to be unnecessary as calibrated concentrations were always in good agreement with concentrations calculated from the weight of Na2S03. 1172 Envlron. Scl. Technol., Vol. 22, No. 10, 1988
Table 11. Summary of Experimental Data initial
no. 1 2 3 4 5 6 7 8 9
after 24 h [S(IV) + [S(IV) + 02, S(IV), S(VI)I, S(IV), S(VI)I, pH mg/L pM rM rM PM Sediment-Water Suspensions (1g/L Sediment) 7.11 8.0 12.8 46 0" 45 7.07 8.0 51 90 22 89 7.07 8.0 225 250 157 245 8.36 8.0 4.0 35 Ob 36 21 82 8.36 8.0 41 83 8.42 8.0 223 272 181 271 0 6.0 33 4.1 36 8.45 8.42 0 41 93 38 90 8.45 0 190 218 169 193
Filtrates (Test Solutions Exposed 2 Days to 1 g/L Sediment Then Filtered before Adding Sulfite) OC 54 10 7.01 7.9 8.3 53 11 7.01 7.9 47 90 26 94 12 7.01 8.0 185 234 143 228 13 8.45 7.9 8.0 49 0.5 52 14 8.45 8.0 88 129 29 125 15 8.45 7.9 270 313 156 290 NaC1-Buffer Solutions (Controls To Which No Sediment Was Added) 16 7.12 8.0 14 24 10 20 88 17 7.12 7.9 99 60 73 15 16 12 18 8.43 8.0 13 82 67 73 19 8.43 7.9 77 331 351 20 8.43 7.9 366 374 462d 21 7.10 0 460 462d 460 450 441d 441d 0 450 22 8.39 Biological Activity Experimentse (1 g/L Sediment Suspensions) 51 75 23 78 23 8.40 8.0 53 71 29 72 24 8.40 8.0 'Reached zero concentration at 23 h. *Reached zero concentration at 19 h. 'Reached zero concentration at 11 h. dMeasured after 48 h. OSamule 23 contains 1% CHCL
All metal analyses were performed with a Perkin-Elmer 2380 atomic absorption spectrometer equipped with a heated graphite furnace atomizer. Secondary standards for calibration were prepared daily, with primary stock solutions prepared by customary procedures (14). The stock solutions were diluted with a solution containing M NaCl so that the standard 0.2% HN03 and 1 X solution matrix was similar to the sample solution matrix. National Bureau of Standards Reference Material 1643-b, "Trace Elements in Water", was analyzed with every batch of samples.
Experimental Results Table I1 summarizes the results. In this table, S(1V) represents SO-: plus HS03-, and S(V1)represents SO:-. The initial sum of S(1V) and S(V1) exceeds initial S(1V) because of sulfate that is leached from the sediment. Typically this was 30 pM sulfate, about one-third the average concentration in river water (15). The values reported 24 h after sulfite addition are based in most cases on interpolation between data points, which were taken over periods of up to 48 h. Examples of complete reaction curves will be presented below. In the following paragraphs we will comment on specific aspects of the experimental results. Role of Oxygen. In the absence of oxygen, little if any sulfite is lost. This is illustrated in Figure 1,which compares runs 5 and 8. This result indicates that the principal loss mechanism is oxidation by 02.Reactions that might occur in the absence of 02,such as reductive dissolution of iron oxyhydroxide and manganese oxyhydroxide grain
-
0.25
0.00 0
10 20 Time (hours)
0.25
1
0.00' 0
"
10
*
'
'
"
"
20 30 40 Time (hours)
50
30
0.75
1
\-
B
1.ool
0.00 0
10 20 T h e (hours)
30
Figure 1. Loss of S(IV) as a function of time without O2(A) and with 0, (6). Data from runs 8 and 5 (see Table I1 for conditions).
coatings on the sediment particles, apparently cause insignificant amounts of S(1V) to be oxidized. As reported later, this has been confirmed by analyzing the solutions for Fe and Mn. Of the three anoxic experiments (runs 7-9, Table 11),only run 9 displays significant sulfite loss over 24 h. Below we will show that this loss is probably due to a nonoxidative process. Role of Particles. The effect of particles on the oxidation of S(1V) by O2can be observed by comparing results for the following three types of solutions: the sedimentwater suspensions, the filtrates from which suspended particles had been removed after 2-day equilibration time, and the control solutions which contained the NaCl and buffer components but which had not been exposed to sediment. Figure 2 presents a comparison of one set of results at pH 7.1 (runs 2, 11, and 17). Sulfur(1V) was oxidized at a much faster rate in the sediment-water suspension and in the sediment-water filtrate than in the control solution. In the control solution, about 15% of the S(1V) disappeared in the first few hours, after which very little loss was observed. In contrast, in both the sediment-water suspension and the sediment-water filtrate, oxidation proceeded a t nearly the same rate, suggesting that the sediment particles themselves had little or no effect on the rate of the reaction. Rather, material released from particle surfaces and capable of passing a 0.2-pm Nuclepore filter must have catalyzed oxidation of sulfite by oxygen. This material includes dissolved trace metals such as Cu, Fe, and Mn (see below), which are known to catalyze this reaction (9, 16). At pH 8.4, as at pH 7.1, S(N)was oxidized faster in both the sediment-water suspension and the sediment-water filtrate than in the control solution (Table 11). However, the reaction in the filtrate was substantially faster than in the suspension in two of the three sets of runs (5 compared to 14 and 6 compared to 15). Apparently the particles themselves have an inhibitory effect in these runs. Oxidation Products and Evidence for Intermediates. The major oxidation product of S(1V) in the sediment-water suspensions and in the sediment-water filtrates was S042-.Both S042- and S2062-have been reported as products of sulfite oxidation, the amounts of each depending on the reaction conditions (9, 16). Our ion chromatographic conditions did not allow determination of S20G2-.Nevertheless, in Table 11, comparison of the initial sum S(IV) + S(V1) with the sum observed after 24 h reveals no systematic change, except in runs 9 and 15
Time (hours)
$+-
:::I 0.00 0
C
,
, 10
,
,
,
,
,
,
20 30 40 Time (hours)
,
, 50
Figure 2. Comparison of S(IV) loss for a 1 g/L suspension of MESS1 (A), a filtrate from such a suspension (B), and a salt/buffer solution that was not exposed to sediment (C). Losses from the suspension and filtrate are much greater than from the control. Data from runs 2, 11, and 17.
and in the NaC1-buffer solutions. These exceptions are discussed below. Conservation of S(1V) + S(V1) implies that no more than a few percent of sulfite is converted to sulfur-containing products, other than sulfate, in these runs. In the NaC1-buffer solutions, which were not exposed to sediment, the small amount of S(1V) lost in 24 h is not replaced by an equivalent amount of sulfate. Yet the loss mechanism definitely involves oxidation; in identical solutions under a nitrogen atmosphere (runs 21 and 22), no S(1V) loss occurred over a span of 48 h. This implies that the oxidation mechanism in these solutions is different than in the solutions exposed to sediment. The most probable cause of this loss of total sulfur is the formation of S2062-, which would not have been detected. Nonoxidative Loss of S(1V). Runs 9 and 15 in Table I1 also exhibited a decreasing trend in the sum of S(1V) + S(V1) larger than can be attributed to analytical error. This is shown in Figure 3. These two runs involved pH 8.4 solutions containing high initial sulfite concentrations. Since run 9 contained no oxygen, it appears that oxygen is not required to account for the loss. This decrease in measured total sulfur was investigated in a special set of experiments to determine if the loss was caused by the formation of an additional product or if the loss was caused by sorption of S(1V) to the sediment. In order to promote sorption concentrated suspensions (20 g/L) were used. Despite the absence of oxygen, an appreciable amount of S(1V) disappeared in the first 2 days. As seen in Table 111, most of the missing S(1V) was not converted to S(VI), because the sum S(1V) S(V1) also decreased markedly. In an attempt to force desorption of sulfite or sulfate, the pH was raised above 11with NaOH, but this did not cause either anion to appear in significant amounts in most attempts. Direct determination of S2062-by the method of Dasgupta et al. (17)indicated that only about 1% of the initial S(1V) was converted to this product. Because no oxygen was present in the system the formation of other sulfur oxyanions, such as S052- (5), is unlikely. In Table I11 samples 6 and 7 are repeats of samples 2 and 5, respectively, except that a greater NaOH concen-
+
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Table 111. Loss of Total Sulfite and Total Sulfite Plus Sulfate in an Anaerobic Suspension Containing 20 g/L Sediment (Concentration in pM) sample
initial pH
bufferd
initial S(1V)
1 2 3 4 5 6c 7'
7.2 7.10 7.61 7.80 8.39 7.1 8.5
HS03-/S032Mops Mops Tris Tris Mops Tris
460 460 460 460 460 400 400
[S(IV) + S(VI)] S(1V) 700 696 708 692 706 900 910
after 2 days plus NaOH, 1dayb S(1V) recovered [S(IV) + S(VI)] S(1V) [S(IV) + S(VI)] pM present
262 250 120 183 129 130 150
505 489 511 463 393 630 650
222 234 94 177 150 130 120
462 470 342 419 396 620 630
300 380
75 92
'Total buffer concentration 1mM, except HSO
