Environ. Sci. Technol. 1990, 24 264-267 I
Effect of Hydrogen Peroxide on the Alkaline Hydrolysis of Carbon Disulfide Scott Elliott
Department of Chemistry, University of California, Irvine, California 927 17 The stability of carbon disulfide in aqueous alkaline solutions has been compared in the absence and presence of hydrogen peroxide, by a headspace gas chromatography technique. Reactant decay is first order in CS2 and OHin solutions containing sodium hydroxide alone. At room temperature, the second-order rate constant 1 X lo9 M-' s-l matches previous measurements for basic hydrolysis and, along with the activation energy 22 kcal mol-', is consistent with rate control in a carbonatelike hydration step. In alkaline hydrogen peroxide solutions, loss is also first order in H20,, suggesting reaction between CS2and the H02- anion with a second-order rate constant in the range 10-100 M-' s-'. Introduction
Carbon disulfide is a raw material in xanthation, and its hydrolysis
CS2 + OH-
kl
products
(1) has received attention both as a secondary step in rayon manufacture (1-3) and as a means for removal from industrial wastewater effluents (4). Hydrolysis of the related species carbonyl sulfide (OCS) has been recognized as the first quantifiable source of the hydrogen sulfides (H2S, SH-, S2-)to the open ocean (5-3, and this development has led directly to the detect.ion of sulfide in open oxic seawater (8,9). CS2 is also a trace seawater constituent (10, 11) and, by analogy with OCS, should be evaluated as an additional sulfide source. A recent major study of the kinetic properties of aqueous carbon disulfide over the pH range 8-1 1included hydrogen peroxide as an oxidizing agent in order to convert hydrogen sulfide products to sulfate (4). CS2loss rates in ref 4 were equal to or greater than those traditionally observed at much higher pH (13-14 in ref 1-3), with characteristic decay lifetimes (e-folding times; t1/,/0.69) of C10000 s. The indication is that H202opens carbon disulfide oxidation and removal channels alternate to hydroxide attack. In the current work, I have investigated this possibility by monitoring CS2 disappearance from solutions covering a wide range of pH and peroxide concentrations. The well-known basic hydrolysis shown as eq 1 is first order in both carbon disulfide and hydroxide ion (12,13),while a faster CS2removal process is first order in hydroxide and H202 and likely involves the peroxide anion HOf as a nucleophile. I conclude that neither hydrolysis of carbon disulfide nor its reaction with HOP- can be competitive with carbonyl sulfide as an open ocean sulfide source. Experimental Section
Reagents. All aqueous solutions were prepared from distilled deionized water saturated with helium for several hours to remove oxygen. In situations with hydroxide ion present in large excess over carbon disulfide (generally pH 12 and above), unbuffered systems were created by dilution of Fischer 10 normal sodium hydroxide. An 0.05 M potassium carbonate/borate/hydroxidemixture from Fischer Scientific was employed as a buffer in some of the pH 10 *Present address: Dept. of Atmospheric Sciences, University of California. Los Angeles, CA 90024. 264
Environ. Sci. Technol., Vol. 24, No. 2, 1990
experiments. Other buffers were provided by Micro Essential Laboratories (Brooklyn, NY) and included a combination of potassium phosphate monobasic and sodium borate at pH 9, a sodium carbonate/bicarbonate mixture at pH 10, and at 11,sodium carbonate/phosphate dibasic. Detailed buffer compositions constitute proprietary information, but can be obtained on an individual basis from the manufacturer. Fischer 30% H202 was diluted to produce the desired hydrogen peroxide concentrations. Liquid carbon disulfide was degassed on a vacuum line before the generation of vapor-phase gas chromatographic calibration curves for the quantification of CS2in the reactor headspace. Vessels. Reactions were carried out in sealed, blackened 300-mL Pyrex round-bottom flasks equipped with ports for half-hole septa. Before each run, the bulbs were acid washed and rinsed several times with the deoxygenated water. In parallel with all the experiments described below, control reactions conducted in the absence of the reagents of interest ensured that CS2is stable with respect to wall losses. In fact, under conditions in the present work, no degradation was apparent over 24-h periods at any pH up to 11. A typical example of carbon disulfide stability is incorporated into Figure 1. In setting up the runs, reactors were filled to capacity with the helium-saturated water, capped, then water removed, and reagents added with a syringe in amounts arranged to leave a 5-mL gas headspace. During these manipulations, occasional injections of helium preserved a slight internal overpressure and minimized the entry of air. Even at the highet temperatures involved in the study (41 "C in the activation energy determinations), the ratio of total CS2 in the liquid to that in the vapor phase was 50, so that the headspace served only as an indicator of losses in solution and could not act as a kinetic source by replenishing the aqueous phase. Most of the runs proceeded at ambient in the laboratory (22.5 f 0.5 "C). Other temperatures were maintained in an oven with thermostat, or in water baths. Procedure. Reactions were initiated with the addition of hydroxide, peroxide, or both. The vessels were shaken at regular intervals and 500-pL samples of the headspace injected onto a gas chromatograph with thermal conductivity detection. Separation of carbon disulfide from products, including carbonyl sulfide and hydrogen sulfide, was achieved on a 5 f t X 0.25 in. stainless steel column packed with Chromosorb 102. The vapor-phase sensitivity limit corresponded to M aqueous CS2. Headspace concentrations were calibrated by comparison with vapor-phase CS2standards prepared fresh twice daily at 0.5, 1.0, 5.0, and 10.0 Torr. Prior to each sample removal, 500 pL of humidified helium was injected into the headspace, again to prevent the entry of oxygen. Except as specified in Table I, the initial carbon disulfide concentration in all experiments was under M. Hovenkamp ( 3 ) showed that side reactions leading to the byproduct trithiocarbonate (CS?-) are restricted to 170 at this level. All CS2 decay curves were tracked over at least 1 full kinetic lifetime (until concentration had decreased by more than a factor of e), and in most cases 2-3, with at least five samples taken per lifetime, and in most cases, well over a dozen total. Rate constants were ex-
0013-936X/90/0924-0264$02.50/0
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0 1990 American Chemical Society
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Time, S x
-5
-4
-3
/
/ I
1
I
-I
-2
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Log,o (concentration), Molar
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Figure 3. Initial carbon dlsulflde removal rates as functions of concentration of CS, ( solid circles), hydroxide ion (open circles), hybogen peroxide at pH 10 (solld triangles), and hydroxide ion at [H202] = lo3 M (open triangles). Data are from Table I. I
I
H202 =
t 0
8
Molar
1
pH = 9
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0.5
L
n
.e
"t
g 0.0 Y
0
IO 20 30 Time, S x 10-3
40
Flgm 2. Sample CS2decay plot for hydrogen peroxide concentration M, in a pH 9 buffer, as tracked by the natural logarithm of headspace pressure.
Table I. Initial Concentrations and Reactions Rates for CS2 Decay with Various Combinations of Reactants, and Rate Constants for Equations 1 and 2 in the Texta
init concn, log,, M CS2 OHpH HzOZ -3.77 -3.26 -2.77 -2.26 -3.77 -3.77 -3.77 -3.77 -3.77 -3.77 -3.77 -3.77 -3.77 -3.77 -3.77 -3.77
-1.0
-1.0 -1.0 -1.0 -1.5 -1.0 -0.5 -0.0 -4.0 -4.0 -4.0 -4.0 -4.0 -5.0 -4.0 -3.0
13.0 13.0 13.0 13.0 12.5 13.0 13.5 14.0 10.0 10.0 10.0
10.0 10.0 9.0 10.0 11.0
-3.36 -2.88 -2.50 -2.36 -2.05 -3.05 -3.05 -3.05
init rate, log10 M/s -7.73 -7.30 -6.86 -6.30 -8.14 -7.76 -7.34 -6.66 -7.53 -7.15 -6.73 -6.58 -6.30 -8.22 -7.04 -6.10
Figure 4. Arrhenlus plot of rate constant measured for reaction 1 as a function of reciprocal temperature (Kelvin).
Table 11. Reaction Order Determined from the Initial Rate Data in Table I
rate, loglo M-1 s-l kl
3.2 3.4 3.6 ( I / T ) x 1000
k2
-2.96 -3.04 -3.09 -3.04 -2.87 -2.99 -3.07 -2.89
CS2
init concn, -log,, M OHH202
3.77-2.26 3.77 3.77 3.77
*
1.0
1.5-0.0 4.0 5.0-3.0
reaction order for species varied
3.36-2.05 3.05
0.94 0.04 (CS,) 0.97 f 0.10 (OH-) 0.96 0.04 (HZO,) 1.06 f 0.07 (OH-)
viations from first-order kinetic behavior. 1.30 1.20 1.24 1.24 1.22 1.30 1.48 1.50
All runs performed at room temperature.
tracted by interpreting CS2 loss over the first e-folding period in pseudo-first-order terms. Linear regression was applied to plots, such as the sample shown in Figure 2, and first-order decay translated into bimolecular rate constants. Hydroxide was either buffered or present in large excess, so that its concentrations remained essentially unchanged. Nevertheless, all vessels were checked for final pH on a pH meter. Hydrogen peroxide levels exceeded those of carbon disulfide but were not always sufficient to be considered invariant. For some of the peroxide experiments, rate constants are based on H202concentrations averaged over 1carbon disulfide lifetime and falling several percent below the initial values in Table I. Within this simple analysis framework, there were no significant de-
Results and Discussion It was readily demonstrated in a series of preliminary orientation experiments that hydrogen peroxide and the hydroxide ion operating together can remove CS2 faster than either by itself. As shown in Figure 1,the disulfide survived on the order of 3-30 h in 0.1 M NaOH or 0.1 M HzOz,but disappeared instantaneously when they were combined. Reaction Order and Rate Constants. In Table I, initial reaction rates are reconstructed for the headspace experiments by extrapolating the measured second-order rate constant values to time zero. A standard initial rates display is given in Figure 3. Slopes for regression of rate against varying reactant concentrations were calculated in the SPSS statistical software package. The results are summarized in Table 11. None of the values differs by more than 6% from unity, and the dashed lines in Figure 3 have been fixed at unit slope to emphasize this point. Earlier studies have implied that uptake of carbon disulfide by pure hydroxide solution is first order in CS2and OH- (e.g., ref 12 and 13). Table I1 verifies the conclusion over wide concentration ranges. The simplest choice for rate-determining step is eq 1leading to the dithiocarbonate species (CS20H-, CS202-;ref 3 and 4). In both this work Environ. Sci. Technol., Vol. 24, No. 2, 1990 265
Table 111. Temperature Dependence of Rate Constant for Hydroxide Attack on CSz
T, "C 41.0 39.0 30.5 22.5 15.0 2.0 2.0 2.0 2.0
init concn, log,, M OHCS2
log,, M-1 s-l
kl,
-1.0 -1.0 -1.0 -1.0 -1.0 -1.0 -1.0
-3.77 -3.77 -3.77 -3.77 -3.77 -3.77 -3.77 -3.77 -3.77
-1.99 -2.12 -2.56 -2.96 -3.42 -4.22 -4.27 -4.20 -4.10
-1.0
-1.0
~~
~
Table IV. Selected Rate Constants and Activation Energy ( E , ) Measurements for Hydration of Carbon Dioxide, Carbonyl Sulfide, and Carbon Disulfide k(298) COZ + HZ0 COz + OH-
+ HzO csz+ H,O CSz + OHOCS
OCS + OH-
0.03 S-' 10' M-' s-l 2X s-I 13 M-' s-'
