Reduction of Acrolein by Elemental Iron: Kinetics, pH Effect, and

Mar 16, 2006 - In this study we investigated the reductive transformation of acrolein by elemental iron and evaluated the feasibility of using iron to...
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Environ. Sci. Technol. 2006, 40, 2765-2770

Reduction of Acrolein by Elemental Iron: Kinetics, pH Effect, and Detoxification SEOK-YOUNG OH, JAEWOO LEE, DANIEL K. CHA, AND PEI C. CHIU* Department of Civil and Environmental Engineering, University of Delaware, Newark, Delaware 19716

Acrolein is a highly toxic R,β-unsaturated aldehyde that is widely used as a biocide, a cross-linking agent, and an intermediate in the chemical industry, among other applications. In this study we investigated the reductive transformation of acrolein by elemental iron and evaluated the feasibility of using iron to detoxify acrolein. At acidic and neutral pH, acrolein was transformed by iron through reduction of the CdC double bond to propionaldehyde. The reduction appeared to involve the chemisorption of acrolein to the iron surface followed by reduction of adsorbed acrolein. Both the adsorption and reduction rate constants decreased with increasing pH. Between pH 7.0 and 7.4, the acrolein adsorption rate constant decreased precipitously, resulting in a sharp decline in its removal rate. At higher pH, acrolein disappeared rapidly in control without iron, presumably due to reversible, base-catalyzed hydration. At equilibrium, approximately 93% of acrolein was hydrated, corresponding to an equilibrium constant of 13. Acrolein at 25 mg/L completely inhibited aerobic respiration; in contrast, its reduction product propionaldehyde was biodegradable. This suggests that elemental iron may be used to pretreat acrolein-containing wastes prior to aerobic biodegradation. To our knowledge, this is the first report of reduction and detoxification of an R,β-unsaturated aldehyde by elemental iron.

Introduction Acrolein (CH2dCHCHO, aka 2-propenal) is a clear, colorless liquid with an acrid odor (1). This R,β-unsaturated aldehyde is used primarily as an intermediate in the production of acrylic acid (1) and also as an aquatic herbicide and microbiocide (2-4). The annual production of acrolein in the United States was 40 000 tons/year in 2003 (5). Because of its rapid binding to cellular materials such as proteins and peptides (1), acrolein is toxic to rats, fish, frogs, bacteria, and humans (1, 6-8). Inhalation of acrolein vapor (∼10 ppm) can be lethal (6). Chronic exposure to acrolein may result in nausea, vomiting, diarrhea, bronchitis, and severe respiratory and ocular irritation (1, 6). The mutagenic and carcinogenic potential of acrolein was also reported (9, 10). Acrolein is on the U.S. EPA’s priority pollutant list and is classified as a possible human carcinogen (class C) (6, 11). It is known that R,β-unsaturated aldehydes can be reductively transformed to saturated alcohol either through unsaturated alcohols by reduction of the CdO double bond or through saturated aldehydes by reduction of the CdC * Corresponding author phone: (302)831-3104; fax: (302)831-3640; e-mail: [email protected]. 10.1021/es052246f CCC: $33.50 Published on Web 03/16/2006

 2006 American Chemical Society

double bond (12, 13). Because unsaturated alcohols are important intermediates for the synthesis of perfumes, flavorings, and pharmaceuticals (12-14), many studies have been conducted to develop processes, including catalytic hydrogenation and electrochemical reduction, to convert R,βunsaturated aldehydes to unsaturated alcohols (e.g., allyl alcohol, CH2dCHCH2OH, in the case of acrolein) (12-22). Hydrogenation of acrolein with catalysts such as Pt, Ru, and Pd preferentially reduces the CdC double bond of acrolein and produces n-propanol through mostly propionaldehyde (or propanal, CH3CH2CHO) (12, 14). Claus and co-workers (13, 18, 19) showed that hydrogenation of acrolein over Au catalysts raised the yield of allyl alcohol by up to 41.5%. The group also showed that hydrogenation of crotonaldehyde (CH3CHdCHCHO) over Ag produced approximately 53% of crotyl alcohol (CH3CHdCHCH2OH) at 473 K (20). Compared to catalytic hydrogenation, electrochemical reduction of acrolein was less examined. Horanyi and Torkos (21) reported that acrolein was transformed to propane through both allyl alcohol and propionaldehyde by electrochemical reduction on a Pt electrode in acidic solution. Results of polarographic reduction of acrolein with limited current at pH 5-7 by Spritzer and Zuman (22) showed that acrolein was reduced to propionaldehyde through electron transfer (eq 1).

CH2dCHCHO + 2e- + 2H+ f CH3CH2CHO

(1)

In addition to reduction, R,β-unsaturated aldehydes can undergo hydration reaction in aqueous solution (23-25). Under highly acidic (pH < 1) or basic (pH > 12) conditions, acrolein was hydrated (eq 2) to 3-hydroxypropanal (or β-hydroxypropanal, CH2OHCH2CHO) (23-25).

CH2dCHCHO + H2O ) CH2OHCH2CHO

(2)

In both acid- and base-catalyzed hydration, approximately 90-92% of acrolein was transformed to 3-hydroxypropanal at equilibrium at room temperature, indicating the hydration rate constant was about 10 times higher than that for the reverse (dehydration) reaction (2, 23, 24). At near-neutral pH, hydration rate of acrolein was found to increase with pH, with a half-life of 50 h at pH 6.6 and 38 h at pH 8.6 (2). In natural waters and soil, acrolein was estimated to have a half-life between 30 and 100 h (2, 4, 26). Smith et al. (3) showed that acrolein was transformed to 3-hydroxypropanal, allyl alcohol, propionic acid, propanol, and 3-hydroxypropionic acid in aerobic and anaerobic sediments over a few days. These studies show that acrolein hydration at neutral pH may be too slow for treatment purposes. In addition, due to its high water solubility (206-270 g/L) (1), the acrolein that remains (8-10%) at equilibrium can still be highly toxic. In recent years, application of elemental iron for water and wastewater treatment was proposed and investigated (27-30). Initially used in permeable reactive barriers (PRBs) for the remediation of chlorinated solvent-contaminated groundwater in mid-1990s, elemental iron has been shown to be effective in degrading many other organic and inorganic pollutants (31). We previously reported that treatment of wastewaters containing azo dyes, nitroaromatic compounds, and heterocyclic nitramines with elemental iron markedly enhanced the rates and extents of mineralization of these refractory compounds in subsequent biological or chemical oxidation processes (28-30). We are conducting two fieldscale studies to demonstrate the feasibility of elemental iron for treating wastewaters containing these compounds. VOL. 40, NO. 8, 2006 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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To date, studies on the reduction of organic pollutants with elemental iron have focused primarily on halogenated and nitrogenous compounds, whereas other types of organic pollutants remain largely uninvestigated. This study was undertaken to (1) assess the reactivity of R,β-unsaturated aldehydes toward elemental iron, using acrolein as a model compound, (2) examine the kinetics of acrolein transformation under different pH conditions, and (3) assess the effect of iron treatment on the toxicity and biodegradability of acrolein.

Materials and Methods Chemicals. Acrolein (90%), propionaldehyde (97%), sodium formate (>99%), sodium acetate (>99%), and 2,4-dinitrophenylhydrazine (DNPH, 97%) were acquired from Aldrich (Milwaukee, WI). MES (2-[N-morpholino]ethanesulfonic acid, >99.5%), MOPSO (3-morpholino-2-hydroxypropanesulfonic acid, >99%), HEPES (N-[2-hydroxyethyl]piperazine-N′-[2ethanesulfonic acid], >99%), TAPS (N-[tris(hydroxymethyl)methyl]-3-aminopropanesulfonic acid, >99.5%), and CAPSO (3-[cyclohexylamino]-2-hydroxy-1-propanesulfonic acid, >99%) were obtained from Sigma (St. Louis, MO). Phosphoric acid (∼85%), acetonitrile (HPLC grade), and methanol (HPLC grade) were purchased from Fisher Scientific (Pittsburgh, PA). The elemental iron used in this study was commercial iron from Master Builders, Inc. (Aurora, OH) and was used without pretreatment. The Master Builders iron was characterized in a previous study (32). The specific surface area of the iron was 1.29 m2/g as measured by the BET method with N2. Batch Experiments. The procedures and conditions for the batch reduction experiments with elemental iron were described in detail in our previous studies (33, 34). Sample preparation and sealing of batch reactors were performed in an anaerobic glovebag (Bell-Art Products, Pequannock, NJ) under N2. Replicate 8-mL borosilicate vials were set up. Each vial contained 5 mL of acrolein solution and 1 g of iron, corresponding to a BET surface area concentration of 258 m2/L. Acrolein solution was prepared by spiking a known amount of acrolein into deionized water that had been deoxygenated by purging with N2 for over 30 min in the glovebag. The initial concentration of acrolein was 0.680 ( 0.013 mM in one set of experiments and 0.587 ( 0.025 mM in another. The solution pH was controlled using the following buffers at 0.1 M unless otherwise noted: formate (0.3 M, pH 4.4(0.5), acetate (0.2 M, pH 5.1(0.4), MES (pH 6.3(0.3), MOPSO (pH 7.0(0.1), HEPES (pH 7.4(0.1), TAPS (pH 8.4(0.1), and CAPSO (9.7(0.1). The significant pH shifts that occurred despite the high buffer concentrations used, especially under acidic conditions, were due to rapid consumption of acidity (proton) during iron corrosion. Vials were shaken at 100 rpm in a horizontal position using an orbital shaker. At different elapsed times, replicate vials were sacrificed and supernatant from each vial was passed through a 0.2-µm cellulose filter (Millipore, MA) for analysis. Control experiments were performed without iron under identical conditions to measure nonredox transformation of acrolein in different buffer solutions. Chemical Analysis. Acrolein and propionaldehyde were analyzed using a Varian HPLC (Walnut Creek, CA) equipped with a SUPELCO LC-18 column (250 × 4.6 mm, 5 µm, Supelco, Bellefonte, PA) and a UV detector (Varian 2510) following derivatization with DNPH (20 mM in acetonitrile, pH 2.5) (35). One milliliter of filtered sample from each vial was derivatized with 1 mL of DNPH solution for 1 h, which was necessary to achieve maximum derivatization efficiency as found in our preliminary tests. To obtain the optimal extraction efficiency for buffered samples, one drop of concentrated phosphoric acid was added to decrease the pH to about 2.5 (36). A methanol-water mixture (65/35, v/v) 2766

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FIGURE 1. Reduction of acrolein with granular iron at pH 4.4 ( 0.5. Fitted curves were obtained using eqs 3-5 (R2 ) 0.973). Error bars are standard deviations based on replicates. was used as the mobile phase at a flow rate of 1.0 mL/min for HPLC analysis. The wavelength for the UV detector was set at 365 nm, and the injection volume for all samples was 10 µL. The LC retention times for acrolein and propionaldehyde were 18.1 and 19.2 min, respectively. The method detection limit for acrolein was 0.003 mM. Biodegradability Assays. Aerobic biodegradability assays were conducted using a Hach BODTrak Apparatus (Hach, Loveland, CO). The apparatus continuously measured change in oxygen partial pressure resulting from aerobic respiration inside a sealed BOD bottle. The system was interfaced with a PC for automatic data acquisition at fixed time intervals. Duplicate 340-mL solutions of acrolein and propionaldehyde (25 mg/L each) were placed in 600-mL dark BOD bottles. Five milliliters of nitrification inhibitor solution (36 g/L, Hach Formula 2533), 20 mL of 0.5 M phosphate buffer, and a BOD nutrient capsule (Hach) were added to each bottle. A BOD seed (Hach) was preaerated and settled, and 5 mL of the supernatant was used to inoculate each sample. Lithium hydroxide powder was placed near the top of each bottle to absorb CO2. The bottles were sealed and placed in a temperature-controlled room at 25 °C for 5 d. Each bottle initially contained about 230 mL of air in the headspace. Cumulative BOD of test samples were corrected against that of the control due to endogenous respiration.

Results and Discussion Under acidic conditions, acrolein disappeared rapidly from solution in the presence of elemental iron. The initial acrolein was removed to below detection within 60 min at pH 4.4 ( 0.5 and in 90 min at pH 5.1 ( 0.4. Concurrently, propionaldehyde was produced in solution almost stoichiometrically (91.6% recovery). A typical result of acrolein transformation with iron at pH 4.4 ( 0.5 is shown in Figure 1. The nearquantitative formation of propionaldehyde indicates the removal was due to reduction, and acrolein was transformed by iron mainly via reduction of the CdC double bond rather than the CdO double bond. Propionaldehyde was stable in the presence of elemental iron for more than 3 h, indicating the carbonyl group was nonreactive compared to the CdC double bond. This is in contrast to catalytic hydrogenation and electrochemical reduction of acrolein (12, 14, 21), where both the CdC and CdO double bonds were reduced, producing n-propanol (12, 14) or propane (21) as end products. Figure 1 also shows that hydration of acrolein at pH 4.4 was relatively slow, as acrolein concentration in controls (without iron) decreased only minimally over the same

period. The slow hydration suggests the reverse (dehydration) reaction can be ignored at pH 4.4, because the rate constant for dehydration is about 10 times smaller than that for acrolein hydration (23, 24) and because the concentration of the hydration product, 3-hydroxypropanal, would be low. If we assume the disappearance of acrolein in the presence of elemental iron was due to both reduction and hydration of acrolein and that these reactions were parallel and first-order in aqueous acrolein concentration, then the disappearance of acrolein and formation of propionaldehyde can be described by eqs 3-5

-

d[acrolein] ) (kR + kH)‚[acrolein] dt

(3)

d[propionaldehyde] ) kR‚[acrolein] dt

(4)

d[acrolein]control d[3-hydroxypropanal] ) ) dt dt kH‚[acrolein]control (5) where kR and kH are first-order rate constants for acrolein reduction and hydration, respectively. kH was estimated through least-squares fitting of control data to eq 5 using Sigma Plot (Systat, Point Richmond, CA), and the kH value was then used to calculate kR by fitting the data to eqs 3 and 4 using Scientist (MicroMath, Salt Lake City, UT). Good fits were obtained for acrolein and propionaldehyde concentrations at pH 4.4 ( 0.5 (Figure 1, R2 ) 0.993) and pH 5.1 ( 0.4 (result not shown, R2 ) 0.985). Note that these results would be essentially the same if hydration had been ignored, as kH was more than 2 orders of magnitude smaller than the corresponding kR at the same pH. At pH 6.3 ( 0.3, acrolein reduction was slower, while its hydration was faster than at pH 4.4 and 5.1. Nonetheless, reduction was still the dominant reaction since 90.2% of acrolein removed was recovered as propionaldehyde (Figure 2a). When eqs 3-5 were used to fit the data, however, a marked difference appeared between observed and predicted propionaldehyde concentrations. There seemed to be an initial lag of approximately 20 min before propionaldehyde was detected in solution. The model also consistently overpredicted propionaldehyde concentrations, in part because the mass recovery at pH 6.3 was lower than that at pH 4.4 and 5.1, especially in early times. These discrepancies suggest that it might not be appropriate to model acrolein reduction by iron as a single-step, first-order reaction. The overprediction of propionaldehyde concentrations and incomplete mass balance suggest that either an intermediate accumulated to a greater extent or adsorption of acrolein or propionaldehyde to iron was more pronounced at pH 6.3. Because reduction of acrolein to propionaldehyde involves the transfer of only two electrons, a one-electron reduction intermediate would be a radical anion, the accumulation of which (as much as 35% of the initial acrolein) would be most unlikely. We showed in this study and a previous study (37) that small saturated aldehydes such as propionaldehyde and formaldehyde did not react with or adsorb to commercial iron. Therefore, the most plausible explanation for the missing mass at pH 6.3 was adsorption of acrolein to iron. Because acrolein is very soluble in water (1), significant hydrophobic adsorption is unlikely. However, it is possible for acrolein to adsorb to iron surface through chemical interactions. This may occur due of the π acidity of the CdC double bond, which may allow acrolein to bind to transition metals having a filled orbital of π symmetry. Alternatively, adsorption of acrolein to iron may occur through electron donor-acceptor interactions, promoted by intramolecular charge separation due to resonance of the conjugated CdC

FIGURE 2. Reduction of acrolein with granular iron at pH 6.3 ( 0.3. Fitted curves were obtained using eqs 3-5 (R2 ) 0.874) for (a) and eqs 5-8 (R2 ) 0.985) for (b). Error bars represent standard deviations.

and CdO double bonds. Indeed, chemisorption of acrolein to metal catalyst surface has been proposed to be responsible for the preferential reduction of the carbon-carbon π bond over the carbonyl group. Through density function theory calculations, Delbecq and co-workers (15-17) suggested that product selectivity in acrolein hydrogenation might be controlled by how the unsaturated aldehyde was adsorbed on catalyst surface. Similarly, Mohr and Claus (13) proposed that the mode of adsorption of unsaturated aldehydes to catalyst might explain the selectivity in their reduction to saturated aldehydes versus unsaturated alcohols. Although the mechanism for acrolein adsorption to elemental iron is not clear, any proposed mechanism must explain the observation that acrolein was transformed through reduction of the CdC double bond, whereas the CdO double bond was inert, in contrast to catalytic hydrogenation and electrochemical reduction. Because reactivity of elemental iron generally decreases with increasing pH (38, 39), the initial lag and pronounced accumulation of adsorbed acrolein at pH 6.3 were probably due to slower reduction of adsorbed acrolein at higher pH. Evidence for acrolein adsorption to iron will be presented later. Acrolein reduction by elemental iron can thus be conceptualized as occurring in four sequential steps: transport of acrolein from solution to iron surface, adsorption of acrolein to iron, reduction of adsorbed acrolein to propionaldehyde, and release of propionaldehyde to solution. Calculations (40) and our preliminary tests indicated external mass transport was not limiting under our experimental conditions. Assuming desorption of propionaldehyde was VOL. 40, NO. 8, 2006 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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rapid, the previous kinetic model can be modified to include adsorption of acrolein, as shown in eqs 5-8

d[acrolein] ) (kH + kA)‚[acrolein] dt

-

(6)

d[acrolein]adsorbed ) dt kA‚[acrolein] - kR‚[acrolein]adsorbed (7) d[propionaldehyde] ) kR‚[acrolein]adsorbed dt

(8)

where kA is the first-order rate constant for acrolein adsorption. Using eqs 5-8, a much better agreement was obtained between observed and fitted propionaldehyde concentrations at pH 6.3 (R2 ) 0.985 for Figure 2b, vs R2 ) 0.874 for Figure 2a). The modified model was able to better explain the initial lag and predicted a maximum acrolein adsorption of 37% at 45 min at pH 6.3. Eqs 5-8 were then used to fit the data at pH 4.4 and 5.1. The result for pH 4.4 is shown in Figure S1 (R2 ) 0.988), for comparison to Figure 1. A good agreement was also obtained for pH 7.0 ( 0.1. Reduction of acrolein by iron was slower at 7.0 ( 0.1: the initial lag increased to 30 min and complete removal took 6 h. However, reduction was still fast relative to hydration, as shown by the final propionaldehyde yield of 82%. The model predicts that the remaining 18% was due to acrolein hydration and adsorption. At pH 7.4 ( 0.1, however, hydration of acrolein was no longer negligible compared to its reduction (Figure 3). Acrolein reacted much more slowly with iron than at lower pH and fully disappeared only after 64 h (3840 min). The initial lag before propionaldehyde formation increased to 120 min. Propionaldehyde accounted for only 41.6 ( 0.5% of initial acrolein after 136 h (8160 min). Over the same time, 87% of acrolein was transformed through hydration in the control. The extent of hydration was approaching the reported equilibrium point (92% (2)). Because of the significant hydration, the reverse reaction; i.e., dehydration of 3-hydroxypropanal to acrolein, needed to be considered. The dehydration reaction was incorporated into the model by modifying eqs 5 and 6

-

d[acrolein]control ) kH‚[acrolein]control dt kD‚[3-hydroxypropanal]control (5′)

d[acrolein] ) (kH + kA)‚[acrolein] dt kD‚[3-hydroxypropanal] (6′)

-

where kD is the first-order rate constant for dehydration of 3-hydroxypropanal. Concentration of 3-hydroxypropanal in control at any time was taken to be the difference between acrolein concentration at that time and initial acrolein concentration. The result of the model fit to the pH 7.4 data is shown in Figure 3. The agreement between experimental data and model fit improved moderately by including dehydration (R2 ) 0.982, vs 0.970 using eqs 5-8). According to the modeling result, 58% of acrolein in the iron reactors was transformed through hydration, whereas 42% reacted through reduction, consistent with the observed propionaldehyde yield. One striking outcome predicted by the revised model (eqs 5′, 6′, 7, and 8) for pH 7.4 was that adsorption of acrolein was minimal throughout the experiment (Figure 3). Based on the fitted rate constants, the marked decrease in acrolein adsorption relative to that at lower pH was due to a sharp decline in kA rather than an increase in kR. This can be seen 2768

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FIGURE 3. Reduction of acrolein with granular iron at pH 7.4 ( 0.1. Fitted curves were obtained using eqs 5′, 6′, 7, and 8 (R2 ) 0.982). Error bars represent standard deviations.

FIGURE 4. Pseudo-first-order rate constants for acrolein reduction, adsorption, and hydration as a function of pH. kH values were obtained from controls without iron, whereas kA and kR were estimated through least-squares fitting using eqs 5′, 6′, 7, and 8. Error bars represent standard deviations. in Figure 4 and Table S1, which summarize the observed and calculated rate constants for adsorption, reduction, and hydration of acrolein at different pH. Figure 4 shows that adsorption became the rate-limiting step for acrolein reduction at pH 7.4, and the sharp decline in kA was largely responsible for the shift in dominant acrolein reaction from reduction (pH < 7.0) to hydration (pH g 7.4). Log kA decreased gradually as pH increased from 4.4 to 7.0 and precipitously as pH increased further to 7.4. Log kR decreased with increasing pH more rapidly than log kA: kR was greater than kA at pH 4.4 and 5.1 and similar to or lower than kA at pH 6.3 and 7.0. This explains why reasonable fits were obtained without considering adsorption at pH 4.4 and 5.1 but not at pH 6.3 and 7.0. Because Figure 4 suggests acrolein adsorption would decrease at higher pH, it might be possible to recover adsorbed acrolein by raising pH. To test this possibility, two replicate vials from the pH 7.4 experiment were sacrificed at 3 and 4 d (4320 and 5760 min) after acrolein had completely disappeared from solution. A 4.5-mL portion of solution was removed from each vial followed by addition of 1 mM NaOH. The vials were immediately shaken for 1 min, and the solution was filtered, derivatized, and analyzed for acrolein. As expected, 45 and 35 nmol of adsorbed acrolein was recovered, corresponding to 0.009 and 0.007 mM at 3 and 4 d, respectively. These concentrations agree well with model

predictions, as shown in Figure 3, and support the acrolein adsorption assumption for the kinetic model. We do not know the reason for the diminished adsorption of acrolein at pH 7.4, although we suspect it was related to changing iron surface charge and/or increasing adsorption of Fe(II) between pH 7.0 and 7.4. We also suspect that adsorption of acrolein to iron, and its reducibility by iron, may be related to the conjugated CdC and CdO double bonds, because alkenes and aldehydes do not appear to adsorb to or react with elemental iron to any noticeable extent. The mechanism for acrolein adsorption to iron may also determine the relative reactivity of the two functional groups, although further research is needed to confirm this. Hydration rate of acrolein increased only slightly as pH increased from 4.4 (2.0 × 10-4 min-1) to 7.4 (2.7 × 10-4 min-1), consistent with Bowmer and Higgins’s result (2). At pH 8.4 ( 0.1, however, hydration completely dominated over reduction: 97% of acrolein disappeared in iron reactors after 64 h versus 93.3% in control over the same time. This was due to both decreased adsorption and faster hydration (1.12 × 10-3 min-1) of acrolein at high pH. No propionaldehyde was detected at pH 8.4. In LC analysis of DNPH-derivatized samples, we observed an unidentified peak with a retention time of 7.8 min that grew as acrolein disappeared. This peak may correspond to the hydration product, 3-hydroxypropanal, although we could not confirm this due to lack of a standard. An attempt to identify 3-hydroxypropanal by extraction with dichloromethane and gas chromatographymass spectrometry was unsuccessful due to poor extraction efficiency and lack of a reference mass spectrum of 3-hydroxypropanal. The kinetics of reversible acrolein hydration at pH 8.4 was described well by eq 5′, as shown in Figure S2. The equilibrium constant was 13.9 based on the final acrolein concentration and 9.7 based on the fitted kH and kD (K ) kH/kD), similar to that (K ) 10) reported elsewhere (25). At pH 9.7, the hydration rate increased further (Figure 4 and Table S1). Addition of elemental iron did not noticeably increase acrolein removal rate, and no propionaldehyde was detected. If one assumes hydration of acrolein was basecatalyzed and kH was first-order in hydroxide concentration, then kH as a function of pH can be described by eq 9. kH2O is the first-order rate constant for reaction of acrolein with water and kOH- is the second-order rate constant for acrolein hydration by hydroxide. The kH2O and kOH- values obtained through least-squares fitting were 2.33 × 10-4 min-1 and 142.4 min-1 M-1, respectively (R2 ) 0.973, Figure 4).

kH ) kH2O + kOH-‚[OH-] ) kH2O + kOH-‚(10pH-14) (9) Biodegradability of Acrolein and Propionaldehyde. Although acrolein can be removed through hydration at high pH, the reaction does not remove acrolein beyond the equilibrium point. In contrast, complete elimination of acrolein was achieved by reduction with elemental iron at neutral and acidic pH. Thus, elemental iron may be used to detoxify wastes containing acrolein and render them treatable by biooxidation processes such as activated sludge. To evaluate the effectiveness of iron treatment in removing acrolein toxicity, aerobic bioassays were carried out for acrolein and propionaldehyde. Figure 5 shows that acrolein was not biodegraded aerobically. In fact, the samples containing acrolein exhibited much lower BOD than the blanks over the 5-d period, indicating aerobic respiration was fully inhibited by acrolein at 25 mg/L (0.45 mM). Note that the toxicity remained despite the fact that a large portion of acrolein was hydrated during the assays (hydration halflife ≈ 1.74 d at pH 7.0). Acrolein was reported to be toxic to organisms (1, 6-8) at concentrations from 0.02 to 2.5 mg/L.

FIGURE 5. Cumulative oxygen consumption during acrolein and propionaldehyde biodegradation. In contrast to acrolein, propionaldehyde was readily degraded. The final BOD of 25 mg/L (0.43 mM) propionaldehyde was 22 mg/L, approximately 40% of the theoretical BOD (without considering growth). While this BOD value suggests propionaldehyde was not fully mineralized by the seed culture, it shows that propionaldehyde was not inhibitory. Thus, toxicity of acrolein can be eliminated through reduction with elemental iron. It may be worthwhile to investigate the feasibility of using iron to detoxify wastes containing R,β-unsaturated aldehydes, such as acrolein and 2-ethyl-3-propylacrolein, to allow treatment of these wastes by aerobic biodegradation.

Acknowledgments We thank GS Engineering & Construction Corporation (Korea) for funding this study.

Supporting Information Available Data of acrolein reduction by elemental iron at pH 4.4 ( 0.5 and fitted curves to eqs 5-8, equilibrium hydration of acrolein at pH 8.4 ( 0.1, and summary of acrolein hydration, adsorption, and reduction rate constants. This material is available free of charge via the Internet at http://pubs.acs.org.

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Received for review November 8, 2005. Revised manuscript received February 13, 2006. Accepted February 20, 2006. ES052246F