Kinetics of Oxygenation of Reduced Sulfur Species in Aqueous Solution Dennis J. O’Brien’* and Francis B. Birkner Department of Civil Engineering, University of Maryland, College Park, Md. 20740
Kinetic studies of the oxygenation of the reduced sulfur species at pH values of 4 , 7 5 5 , and 10 were made. The reaction was determined to be first order with respect to reduced sulfur a t all p H values and nearly first order with respect to oxygen a t pH 7.55. Reaction products were determined, and an overall kinetic model was developed to more fully understand the complex reaction mechanism involved. The model suggested parallel reactions between reduced sulfur species and oxygen to form sulfite, thiosulfate, and sulfate. Reasonable predictions of reactant and product concentrations were obtained.
The occurrence of reduced sulfur species in natural waters is caused either by natural geochemical and biological processes or by the discharge of sulfur-containingwastes into the environment. Many natural waters may contain total reduced sulfur concentrations as high as M; however, occasionally higher concentrations are found in hot spring waters, eutrophic lakes, or in waters that are subjected to certain types of industrial waste pollution. Perhaps the most important natural sources of reduced sulfur species in the aqueous environment result from the biological reduction of sulfates by sulfate-reducing bacteria and from the putrefaction of sulfur-containing amino acids. The discharge of certain types of industrial wastes, for example, acid mine drainage, petroleum wastes, leather tanning wastes, and certain textile wastes, may also be responsible for high localized concentrations of reduced sulfur species in the aqueous environment. Effects of reduced sulfur species in natural waters include the disruption of biological waste treatment processes, an increased chlorine demand, and corrosion of sewers due to acidity resulting from oxidation of sulfides. Finally, low concentrations of reduced sulfur in water supplies or waste streams are aesthetically undesirable and toxic to many aquatic organisms. T o evaluate the economic and ecologic importance of the presence of reduced sulfur species in aqueous environments, it becomes necessary to first understand the mechanisms by which reduced sulfur species are removed from natural waters. Reduced forms of sulfur may be removed from natural waters by chemical precipitation with transition or heavy metal ions, volatilization of the H2S species in aeration processes, biological oxidation, and chemical oxidation by molecular oxygen or other oxidants. This paper will be concerned only with chemical oxidation as a removal mechanism.
Chemistry of S ( -11) Species Hydrogen sulfide is a weak diprotic acid and therefore can exist in three chemical forms in solution, HsS(aq), HS-, and S2-. The solution pH determines the distribution of these reduced sulfur species and is the predominant solution variable. Other parameters that have important effects are ionic strength, buffer capacity, and temperature. Acidity constants for hydrogen sulfide are 7.0 and approximately 14. Thus, in the pH range of natural waters, the bisulfide ion (HS-) is the primary species, while the hydrogen sulfide molecule [HzS(aq)] becomes the predominant reduced sulfur species
Present address, Department of Civil Engineering, University of New Hampshire, Durham, N.H. 03824. 11 14
Environmental Science & Technology
below pH 7. A knowledge of the distribution of reduced sulfur species is important in the interpretation of kinetic behavior. Polysulfide species are formed when neutral sulfur atoms combine with monosulfide species. They can be represented by the chemical formulas, HpS,, HS;, and S:-,where x = 2 - 5. Equilibrium calculations show that the tetrasulfide and pentasulfide species should be the predominant polysulfide forms in neutral and slightly alkaline solutions, but recent experimental work ( I ) has detected only the pentasulfide species. In previous papers, Cline and Richards ( 2 ) and Chen and Morris ( 3 )studied several aspects of the S(-II)T-oxygen reaction. Cline and Richards measured rate constants for the decrease in S(-II)T and 0 2 concentrations as well as product concentrations for the sulfide oxygenation reaction in seawater a t 9.8 “C. In a more comprehensive investigation, Chen and Morris described the complex nature of the dependence of the reaction rate on pH a t intermediate concentration levels. I t was postulated that the reaction proceeds through a mechanism involving polysulfides as intermediates. The present study was undertaken to study the kinetics of reduced sulfur oxidation by molecular oxygen in simulated natural water environments. The method used was to examine the products of the reaction and their rates of formation and change to discern a reaction model from the complex set of interactions in the S(-11)-On-products mixture. With the utilization of standard kinetic techniques, the data for the buildup of the various product species, i.e., sulfite, thiosulfate, and sulfate, were analyzed to describe the reaction by a series of parallel and consecutive reaction rate expressions. In this manner, an overall model was found that was used to predict the product concentrations for various reaction times.
Experimental Standard spectrophotometric techniques were employed to determine the concentrations of the sulfur oxyanions in the reaction mixture. These methods were checked for interferences from the reactants, oxidation products, and the buffers. Total reduced sulfur, S(-11)~, was determined by the amine-sulfuric acid method of Budd and Bewick ( 4 ) .The method is based on the oxidation and transformation of p aminodimethylaniline by ferric ion and hydrogen sulfide to methylene blue. Maximum absorbance of the methylene blue occurs at 670 mu, and Beer’s law is obeyed in the 1.5(10)-6 to 8(10)-s M concentration range. Urban’s spectrophotometric method ( 5 ) ,which involves the cyanolysis of the thiosulfate species and subsequent complexation by ferric ion, was used for the thiosulfate determination. The ferric thiocyanate complex absorbs a t 460 mu, and concentrations as low as M can be determined. Reduced sulfur species caused a positive interference in this analysis but were removed by stripping with nitrogen a t low pH. Sulfite was measured by the West and Gaeke colorimetric procedure (6). The method was modified to prevent sulfide precipitation by elimination of the mercury complexing agent. Concentrations as low as 8 X lo-$ M could be determined. Oxygen concentrations were measured with a Fisher Model 29 gas partitioner that was modified according to Swinnerton et al. (7) to allow the use of large liquid samples.
Qualitatively, polysulfides can be detected by the appearance of a yellow-green color. Quantitative methods for polysulfides are complicated by the high reactivity of these compounds with oxygen and other oxidants and the complex equilibria that exist between polysulfides and monosulfides in solution. In a recent article, Chen and Gupta (8) proposed a method for computing polysulfide concentrations based upon the use of molar extinction coefficients measured by Schwarzenbach and Fisher (9). However, due to the uncertainty about which polysulfide species are actually present in aqueous solution, this method was not used in this study. Instead, solution absorbances a t 290 mu were used as a qualitative indication of the presence of polysulfides. Extensive spectrophotometric analysis of reaction solutions a t 290 mu failed to detect the presence of polysulfide species. Colloidal sulfur in the reaction system could easily be detected visually since its solubility is 5(10)-6 M (10). Its appearance was never detected in any experiments reported herein. Also, no satisfactory method for the determination of sulfate a t low concentrations in the reactor solution was found. Sulfate concentrations were calculated by difference under conditions where colloidal sulfur was not observed and polysulfide absorbance was negligible. Buffers. Experiments presented herein were conducted a t three pH values, 4, 7.55, and 10, and were buffered with acetic acid-acetate, tris(2-amino-2-(hydroxymethyl)-1,3propanediol), and carbonate-bicarbonate buffer systems, respectively. A series of experiments was conducted in sealed bottles to determine the effect of p H on the distribution of oxidation products. These experiments were carried out in a carbonate buffer system a t pH values between 8 and 11. The solution ionic strength was maintained constant in the product distribution runs a t 0.155 M. Reaction Vessel. The reaction vessel used in this study was similar in design to the one employed by Cline and Richards ( 2 ) .The reactor had a capacity of 1.8 L and was equipped with a movable piston which could displace Y2 L of fluid. This made it possible to study the reaction kinetics in a homogeneous system without the presence of a gas phase. Constant temperature water was circulated through the reactor jacket with a centrifugal pump for temperature control, and a pool of distilled water was placed around the piston to provide an airtight seal. Stirring was provided by the use of a magnetic stirrer. The temperature in all experiments was 25 f 0.2 "C. Experimental Procedures. As pointed out by many investigators (2, 3, 11, 12), the oxygenation of S(-II)T species is affected by small quantities of transition metals and organic compounds. Therefore, great care was taken in acid cleaning and washing all experimental equipment. T o minimize the effects of biological catalysis on the kinetic studies, all buffer solutions were autoclaved prior to use. The reaction vessel was then washed with 1/1 (v/v) HC1 to remove any traces of metal ions or sulfur oxidation products that may have been present on the surfaces of the reaction vessel, and the vessel was rinsed thoroughly with sterilized distilled water. The buffer solutions were then oxygenated to near saturation with a pure oxygen stream, and the reactor was sealed. The oxygen concentration was monitored over a period of 18-24 h to ensure that it had stabilized a t a constant value. Prior to the initiation of each experiment, the oxygen concentration, pH, and temperature of the reaction solution were checked. The reaction was initiated by injecting an aliquot of a fresh stock sulfide solution into the reaction vessel. The stock sulfide solution was prepared by dissolving prewashed crystals of sodium sulfide nonahydrate (Na2SSH20) in deoxygenated distilled water. Aliquots for total reduced sulfur, sulfite, and thiosulfate analyses were withdrawn a t periodic intervals from the reac-
tion vessel by advancing the displacement piston into the reactor and forcing the samples into pipets connected to the sampling port. Samples for oxygen analyses were withdrawn with a hypodermic syringe. Upon injection of the reduced sulfur aliquot, the reaction vessel was stirred for 5 min, after which the initial samples were taken. A t the conclusion of the experiment, the pH was measured to determine if it had changed over the course of the reaction. All chemicals used in the preparation of reaction solutions and reagents for the spectrophotometric analyses were reagent grade and were used as received. Stock S(-II)T solutions were prepared with great care to exclude oxygen and were used immediately, since the oxidation products of reduced sulfur species can retard the sulfide oxygenation rate (13) and change the distribution to reaction products (12).
Res u 1t s O r d e r w i t h Respect t o Reduced Sulfur. The reaction order with respect to total reduced sulfur was determined a t pH values of 4, 7.55, and 10.At p H 4 and 10 the order was obtained by following the decrease in the S(-II)T concentration with time under conditions where oxygen was in excess. In the low concentration range, 10-4-10-5 M, the reaction behaved according to a semilogarthmic first order rate law (Figures 1 and 2). Reaction rate constants were calculated to be 0.524 M-l min-l and 2.24 M-l min-' a t pH 4 and 10, respectively. At pH 7.55 the reaction order was evaluated by a differential analysis of the kinetic data, utilizing the method of initial rates.
cy and 13 are the reaction orders with respect to reduced sulfur and oxygen, respectively; k is the reaction rate constant, and the brackets indicate molar concentrations. Written in logarithmic form, Equation 1becomes:
log R, = log k
+ a log [S(-II)T]O+ 13 log [ 0 2 ] 0
(2)
From Equation 2, a plot of log Ri vs. log [s(-II)T]O will have a slope of N , the order with iespect to reduced sulfur. Initial rates of reaction were experimentally determined a t initial reduced sulfur concentrations ranging from 2.2 X 10-5 M to 1.21 X M, and the plot is shown in Figure 3. A linear regression analysis yielded a value of 1.02 f 0.08 for the order with respect to total reduced sulfur. I
pH=4
[02]= l.04X 10%
10.0-
181.75 X IO-% T=25OC
-
9.0 8.0 70-
a
e
-
=
60-
0
5.0-
"
+ -
4.0
W
i
2.0.
1.01 0
I
100
200
1
300
I
400
I
500
I
1
600
TIME,MINUTES
Figure 1. First order 4
plot, reduced sulfur concentration vs. time at pH Volume 11, Number 12, November 1977
1115
Order with Respect to Oxygen. The order with respect to oxygen was determined a t p H 7.55 and a reduced sulfur concentration of approximately M by the initial rate method. As before:
and in logarithmic form: log Ri = log k
+
(Y
log [s(-II)T]O
+ p log [ 0 2 ] 0
(4)
The order with respect to oxygen, /3, is obtained from a plot of log Ri vs. log [O& (Figure 4). For initial oxygen concentrations of 0.208 - 1.07 X M, the oxygen order was 0.80 f 0.25. The overall reaction rate expression for the oxygenation of reduced sulfur species under the stated conditions can thus be written as:
I
I
1
I
1
I
4.0
pH*IO
[Od =I.OXIO-3M I =1.75 X10-2Y
2.0
Ts25OC
I
.I
I
0
100
I
200
I
1
I
1
500
300 400 TIME, MINUTES
600
Based upon the experimental results presented here, subsequent kinetic modeling for the purpose of elucidating an overall reaction model utilized simple reaction rate expressions with reaction orders of unity for oxygen and total reduced sulfur. Product Distribution and Reaction Models. An objective of this study was to formulate an overall model to describe the oxidation of reduced sulfur species under a specific set of solution conditions by measuring the concentration-time dependence of the reactants and products such as sulfite and thiosulfate. Reaction rate expressions for simple parallel and consecutive formation reactions of these products were combined in an attempt to describe concentration profiles of the reaction products, sulfite, thiosulfate, and sulfate. The validity of the particular model chosen was evaluated by comparison of the model predictions with the actual concentration data. The proposed reactions are overall formation reactions that are not necessarily indicative of the actual molecular combinations that occur. Five product distribution experiments were conducted in which total reduced sulfur, oxygen, sulfite, and thiosulfate concentrations were measured with time. Qualitative observations for sulfur and polysulfide were also made. Rate constants for the decrease in reduced sulfur and for the increase in sulfite and thiosulfate were determined by the initial rate method. This information is summarized in Table I. Since sulfate was not experimentally measured, no rate constant for its formation was calculated. Once the rate constants were available, many different combinations of simple bimolecular reactions were tested. Parallel and consecutive reaction sequences were investigated. Reaction models were evaluated by comparison of modelpredicted and actual concentrations for total reduced sulfur and the reaction products. Model Formulation. Rate constants for the production of sulfite and thiosulfate were determined by the initial rate method.
J
Figure 2. First order plot, reduced sulfur concentration vs. time at pH 10
-
/I
[0,]=IO-3M pH= 7.55
.-C E.
f. P
I =.ISSM T -25OC
5
$*.eo*
/
.25
4-
X
F: 3 -
&G 1.
::
U = I . O 2 i .08 pH 7,55 I=.155M T = 25OC
W
0.5K
I
.I
.5
I
1.0
I
!
,
I
203.04,05.0 [S(-II)T]o XIO?M
I
10.0
I
20.0
I
Figure 3. Effect of initial reduced sulfur concentration on initial rate of reaction CY
is order with respect to total reduced sulfur
1116
Environmental Science & Technology
I
I
I
I
I
I
I
I
I
I
!
I
Flgure 4. Effect of initial oxygen concentration on initial rate of reaction
p
is order with respect to oxygen
The initial rates were experimentally determined in the early stages of the reaction, and the rate constants were calculated from Equations 6 and 7. To proceed with the reaction modeling, it was necessary to describe the oxygen concentration with time. Since the order with respect to oxygen was approximately unity, an exponential decay expression was satisfactorily fit to the oxygen data for all reaction times. [021 =
[02loe-rt
(8)
y, the decay constant, is experimentally determined from oxygen data. The value used in this study was 1.91 f 0.14 X
min-l (Table I). The oxidation of sulfite was also studied. In the presence of oxygen, sulfite was oxidized to sulfate (14). It is experimentally impossible to determine the rate constant for this reaction when S(-II)T species are present because sulfite would be produced as a result of the reaction between reduced sulfur species and oxygen, thus preventing an accurate accounting of the change in the sulfite concentration. Therefore, the value of the rate constant for the reaction between sulfite and oxygen was determined by applying a steady state hypothesis to the rate of change of sulfite with time. The concentration of sulfite changed very little from 8 to 24 h. Consequently, after a reaction period of 24 h, it was assumed that the sulfite species attained a steady state concentration, i.e., the rate of sulfite formation from the oxidation of S(-II)T is equal to the rate a t which sulfite is oxidized to sulfate. With this assumption, the rate constant for the oxidation of sulfite, k3, can be determined from experimental and kinetic data. The steady state hypothesis can be represented by:
Table 1. Summary of Kinetic Data for Oxygenation of Reduced Sulfur Species at pH 7.55 lnltlal concn,
M Expt
S(-II)T
1 2 3 4
5
1.10 1.18 1.04 1.17 1.21
x
Rate constants, M-I mln-1
103 02
k
kl
k2
k3
Y
1.02 0.93 1.07 0.85 1.21
1.12 1.37 1.75 1.97 0.97
0.226 0.331 0.553 0.521 0.207
0.110 0.204 0.411 0.220 0.093
1.56 1.25 1.42 2.51 1.26
1.96 2.06 1.95 1.70 1.86
Rate Expressions
The rate expressions can be integrated simultaneously with the aid of a digital computer to yield predicted concentration profiles for all reactant and product species. Comparisons of predicted concentrations with measured quantities for the five product distribution experiments are presented in Figures 5-9. Sulfate data are not presented since sulfate concentrations were not measured. Discussion
(12) In the above analysis it is assumed that the rate of oxidation of sulfite is first order with respect to both sulfite and oxygen. Values of k 3 for five experiments are summarized in Table I. Thiosulfate can be considered a stable reaction product. Little oxidation occurred over 72 h (14), a finding confirmed by Avrahami and Golding (12). The thiosulfate concentration could best be predicted with a rate expression which is zero order with respect to oxygen. This suggests that the mechanisms for thiosulfate and sulfite formation are quite different. The overall reaction model that was developed in this study to explain the behavior of reactants and products in the reduced sulfur oxygen system can be summarized as follows Reactions
Under the experimental conditions of this work, the reaction between reduced sulfur and oxygen is first order with respect to reduced sulfur and approximately second order overall. Kinetically, it is reasonable to assume that the rate of reaction between oxygen and reduced sulfur species is controlled by a bimolecular collision between the reacting species. Reactions in the liquid state are highly unlikely to involve more than two molecules per collision (15). If, as several investigators (2,16,17) suggest, the oxidation proceeds via a free radical chain mechanism following an initiation step consisting of a combination of the HS- molecule and oxygen, the succeeding free radical reactions would presumably be very rapid and the rate of reaction would be controlled by the slower 02-HS- step. Previous work by Chen and Morris as well as Cline and Richards has led to similar overall reaction orders. Chen and Morris ( 3 ) reported an overall order of 1.9 with individual reaction orders 1.34 and 0.56 for S(-II)T and oxygen, respectively. Cline and Richards fitted their data to a second order overall model in which each reactant followed first order behavior. The experimental conditions for these investigations were different from those reported here, and with the existing lack of knowledge of the mechanism of the reaction, little can be concluded concerning the similarities and differences in the results. Several experiments were conducted to determine the effect of solution ionic strength on the rate of the oxygenation reaction. Table I1 summarizes these results. The effect of ionic strength is substantial, a finding also reported by Alferova and Titova (18). This effect can help to explain some widely varying results in different investigations. The results from the ionic strength experiments provide further evidence for a mechanism involving a rate-controlling step between two bisulfide-oxygen ionic species, as originally proposed by Abel Volume 11, Number 12, November 1977
1117
(16).The primary salt effect predicts that for ionic reactions between ions of similar charge sign, an increase in ionic strength will lead to an increase in reaction rate. Thus, Abel postulated an autooxidation mechanism based upon a species such as HSOT. Frost and Pearson (19) define autooxidation as the slow reaction of molecular oxygen with most organic and inorganic substances in the liquid state at moderate temperatures. In Abel's hypothesis, bisulfide and oxygen would first combine to form an activated complex after which various free radical or other reactions would yield the observable reaction products. The results of this study are consistent with Abel's hypothesis.
Reaction Products A summary of the reaction products that have been reported in the literature and the conditions under which they were formed is presented in Table 111. The most commonly observed species are sulfite, thiosulfate, and sulfate. Sulfur (SO)has been detected only occasionally by Avrahami and Golding (12),while Chen and Morris observed its formation only at high [s(-II)T]O/[02]0 ratios in mildly alkaline solution. Several factors can explain the disparity in product distributions observed in Table 111.The solution pH in this reaction system determines the distribution of reduced sulfur species. From analysis of reactor solutions in which H2S is the predominant form of reduced sulfur, the major reaction product was assumed to be sulfate [measured by difference ( 1 4 ) ] .In
solutions where the HS- species predominates, sulfite, thiosulfate, and sulfate are formed as reaction products. Furthermore, Table IV indicates that the product ratios do not change appreciably over the pH range 7.55-10.7. Neither sulfur nor polysulfides were observed in the reaction mixture during the course of this investigation. Thus, it appears that solution pH affects the distribution of reaction products through the redistribution of reactants, primarily H2S and HS-, via simple acid-base equilibria. Apparently, each form of reduced sulfur is oxidized via a different pathway which leads to a different distribution of reaction products. Another reaction parameter that seems to have some effect on the distribution of reaction products in the reduced sulfur-oxygen system is the initial [s(-II)T]/[o]2 ratio. Results from this study are in agreement with those of Cline and Richards ( 2 ) and suggest that as oxygen becomes the limiting reactant, the initial production of sulfite increases. Chen and Morris ( 3 ) and Demirjian ( I I ) , concluded that sulfur precipitation was favored by a high [s(-II)T]/[o2] ratio while the production of sulfite, thiosulfate, and sulfate occurs in systems where the [s(-II)T]/[o2] ratios are small. Another prerequisite for sulfur formation seems to be a high total reduced sulfur concentration (3, 17). To summarize, the [s(-II)T]/[o2] ratio affects the distribution of reaction products in the following ways: A high ratio, along with a total reduced sulfur concentration greater than M, leads to the formation of sulfur.
i 2.4
12.01
12.0I
12.4
- 2.0
e
*-
z
6,0-
A,SOC2 DATA o,S20;2 DATA
g
6L
O , S ( - E I T DATA
4.0- --,PREDICTED
- 1.2
x
-0.8
8
t-
a CONCENTRATIONS
3
A
'0
200
100 TIME
300
400
500
(MINUTES)
TIME (MINUTES)
Figure 5. Model predictions vs. experimental data, experiment 1. pH
= 7.55
Figure 7. Model predictions vs. experimental data, experiment 3. pH = 7.55
L.
r.
0 6,0- o,s(-niT DATA AS0;'DATA
i-
U
5
.s,o;2
-
0-
2 X 1.2 4:
6,0-
t-
0
DATA
4,O---.PREDICTED
CONCENTRATIONS
0,s (-n)TDATA A,S0i2DATA
- 1.2
X
n 0
a Y
A-0.8
k
Y
TIME ( M I N U T E S )
Figure 6. Model predictions vs. experimental data, experiment 2. pH
= 7.55
1118
Environmental Science & Technology
TIME
(MINUTES)
Figure 8. Model predictions vs. experimental data, experiment 4. pH
= 7.55
A low ratio favors the production of sulfite, thiosulfate, and sulfate. The distribution of these reaction products, however, is not affected by solution pH over the range 7.511. At low pH (below pH 6) in dilute solution, the favored reaction product is presumably sulfate. Above pH 7 and for M, sulfite, thiosulfate, and S(- 1 1 ) ~ concentrations below sulfate are significant products. Results from this study agree quite well with those of Cline and Richards ( 2 ) who observed that the reaction products, expressed as percent of the total initial amount of reduced sulfur, were distributed in the following manner: 30-35% as thiosulfate, 10-15% as sulfite, and the remainder presumably as sulfate. Reaction Model One purpose of this investigation was to develop an overall model to describe the oxygenation of reduced sulfur species which would be in accordance with the observed distribution of reaction products. The approach used was to assemble simple rate expressions of parallel and series reactions after experimentally obtaining the needed rate constants.
2.4
10.0-
80
-
t
0
x 6.0 9
5 d
O , S ( - I I ) ~DATA - A,SD;' DATA 0 IS203-2 DATA
4.0- --,PREDICTED
CONCENTRATIONS Y
4.0-0.2 s203-2
' 0
100
200
300
400
500
TIME (MINUTES1
Flgure 9. Model predictions vs. experimental data, experiment 5. pH = 7.55
Table II. Effect of Ionic Strength on Oxygenation of Reduced Sulfur Species * at pH 7.55 Reaction rate constant, k , M-' rnln:'
Ionic strength, molIL
a
0.155 0.97-1.97 1.78 3.97 Initial reduced sulfur and oxygen concentrations were IOT3 M.
The model that best describes the experimental results involves the reaction of reduced sulfur and oxygen in parallel reactions to form sulfite, thiosulfate, and sulfate. This property of the model was demonstrated from several observations. First, the concentrations of sulfite, thiosulfate, and sulfate built up in a steady progressive manner with no discernible lag periods which would have indicated the presence of a rate-controlling intermediate step. Also, product concentrations leveled off concurrently with the reduction in reduced sulfur oxidation rate. Stoichiometric calculations for four experiments are presented in Table V. At a reaction time of 24 h the molar ratio of oxygen consumed to reduced sulfur oxidized, A [ 0 2 ] / A[s(-II)T], was experimentally determined. This can be compared with the following hypothetical stoichiometric reaction: 4HS-
+ 5 -21 0 2 ---+S20T2 + SOB2 + SOT2 + 2H+ + H2O
(21) The theoretical A[o2]/a[s(-II)T] for this reaction is 1.38. The experimental average value of 1.36 further suggests that the oxidation of reduced sulfur yields equimolar concentrations of sulfite, thiosulfate (as S),and sulfate. The results of the model predictions are presented in Figures 5-9. The experimental data are in good agreement with the model predictions up to a reaction time of 500 min, although the model consistently underestimates the sulfite concentration. In this study the initial rate of formation of the products significantly affects the final product distribution. The product distribution after 24 h changes very little as long as no additional oxygen is introduced into the system. The agreement between the model proposed in this study and the observed distribution of reaction products does not provide conclusive proof that the reaction pathways of the overall model actually describe the series of chemical reactions that occur. Any proposed model must also satisfy thermodynamic and energetic considerations. Observations regarding the effects of solution conditions and experimental parameters should also be compatible with proven chemical theories. The results of this investigation, although they were unable to provide a detailed chemical mechanism for the oxygenation of reduced sulfur species, have been shown to describe the overall behavior of the reduced sulfur-oxygen system with regard to the above requirements. In summary, the following general conclusions may be advanced: The reaction between reduced sulfur species and oxygen is first order with respect to both reduced sulfur and oxygen under the reaction conditions of this study. The distribution of reaction products is adequately described by an overall model consisting of a series of parallel reactions to form the major reaction products, sulfite, thio-
Table 111. Summary of Reaction Products Observed in Investigations of Oxygenation of Reduced Sulfur Species Investigator
Chen 8 Morris (3) Avrahami & Golding ( 72) Cline & Richards ( 2) Skopintsev et al.
PH
Reaction solutlon
[S(-11)r10/[0210
Products obsd
6-12 11-14
Controlled Controlled
0.06-1.25 0.08-0.67
si2,
7.8
Seawater
0.125-0.5
so;',
s2032, so42
8.2
Seawater
0.2-8.0
so,',
s20,'
7, 8.6 9-13
Controlled Controlled
0.03-5.0 20
so, soy2,s2032, SOT2 SO32,
s2032,
4-10.7
Controlled
1.0-1.37
so,2,
s2032, SOT2
SO
so, s20,*, soy2 (occasionally), s20i2,SO;'
(20) Demirjian ( 7 7) Alferova & Titova (78) This work
SOT2
Volume 11, Number 12, November 1977
1119
Table IV. Effect of pH on Reaction Product Distribution Product dlstrlbutlon at 24 h [s203-21/4[s( -11 111 0.162 0.169 0.194 0.180 0.223 0.207 0.178 0.230 a Initial concentrations of total r e d d sulfur and oxygen were M. The product disvibuticm is represented by the ratio of the concentration of the product formed to the amount of reduced sulfur oxidized.
PH
[ s o 3 - 2 ] ~ ~ ~ ( - ~ ~ h l
8.1 9.3 9.8 10.7
Table V. Stoichiometry of Reduced Sulfur Oxygenation at pH 7.55 Ex@
Mol ratlo of oxygen consumed to reduced sulfur oxldlzed at 24 h, a[OZ]/a[s(-II)T]
1 2 3 4
5
Average
1.61 1.17 1.33 1.07 1.60 1.36 f 0.25
sulfate, and sulfate. Further conversion of sulfite to sulfate occurs slowly. Caution should be urged in the application of these results. The catalytic effects of many metal ions are well known, and many organic compounds may also affect the reaction rate. The results are useful for a quantitative estimation of the uncatalyzed chemical oxidation rate and for a further understanding of the nature of the complex reaction between reduced sulfur species and oxygen.
Nomenclature I = ionic strength, m o l L k = second order rate constant for the oxygenation of reduced sulfur, M-l min-l Izl = rate constant for the production of sulfite ion, M-l min-1 k z = rate constant for the production of thiosulfate ion, M-’ min-l k s = rate constant for the conversion of sulfite to sulfate, M-’ and min-l
1120
Environmental Science 8 Technology
M = mo1L
Ri = initialrate S(-II)T = sum of sulfur species in the -11 oxidation state T = temperature Aoz/A[s( = stoichiometric ratio, moles of oxygen consumed per mole of reduced sulfur oxidized [SO,*]/A[S( -II)T] = sulfite formation ratio, moles of sulfite formed per mole of reduced sulfur oxidized [SZO,~]/A[~(-II)T]= thiosulfate formation ratio, moles of thiosulfate formed per mole of reduced sulfur oxidized a = reaction order with respect to reduced sulfur /3 = reaction order with respect to oxygen y = oxygen decay constant, min-l
Literature Cited (1) Pringle, D. L., PhD dissertation, Iowa State University, Ames, Iowa, 1967. (2) Cline, J. D., Richards, F. A., Enuiron. Sci. Technol., 3 (9),838-43 (1969). (3) Chen, K. Y., Morris, J. C., ibid., 6 (6),529-37 (1972). (4) Budd, M. S., Bewick, H. A., Anal. Chem., 24,1536-8 (1952). (5) Urban, P. J., 2. Anal. Chem., 179,415-26 (1961). (6) West, P. W., Gaeke, G. C., Anal. Chem., 28,1816-19 (1956). (7) Swinnerton, J. W., Linnebom, V. J., Cheek, C. H., ibid., 34,483-5 (1962). (8) Chen, K. Y., Gupta, S. K., Enuiron. Lett., 4 (3), 187-200 (1973). (9) Schwarzenbach, G., Fisher, A., Helu. Chim. Acta, 169, 1365-90 (1960). (10) LaMer, V. K., Kenyon, A. S., J. Colloid Sci., 2,257-64 (1947). (11) Demirjian, Y. A., PhD dissertation, University of Michigan, Ann Arbor, Mich., 1970. (12) Avrahami, M., Golding, R. M., J. Chem. SOC.,A, 1968, p p 647-51. (13) Chen, K. Y., Morris, J. C., paper presented a t 5th Int. Conf. on Water Pollution Res., San Francisco, Calif., July-August 1970. (14) O’Brien, D. J., PhD dissertation, University of Maryland, College Park, Md., 1974. (15) Amdur, J., Hammes, G., “Chemical Kinetics-Principles and Selected Topics”, p 109, McGraw-Hill, New York, N.Y., 1966. (16) Abel, E., Monatsh. Chem., 87,498-502 (1956). (17) Bowers, J. W., Fuller, M. J., Packer, J. E., Chem. Ind., 1966, p p 65-6. (18) Alferova, L. A,, Titova, G. A., Zh. Prikl. Khim., 42 (l),192-6 (1969). (19) Frost, A. A., Pearson, R. G., “Kinetics and Mechanism”, 2nd ed., p 45, Wiley, New York, N.Y., 1961. (20) Skopintsev, B. A., Karpov, A. V., Vershinena, D. A., Sou. Oceanogr., No. 4, 55 (1964).
Received for review October 26,1976. Accepted June 28,1977. Presented at the Division of Environmental Chemistry, 169th Meeting, ACS, Philadelphia, Pa., April 1975. Work supported in part by EPA Research Fellowship to D.J.O.
