Carbon Tetrachloride Degradation by Alkaline Ascorbic Acid Solution

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Carbon Tetrachloride Degradation by Alkaline Ascorbic Acid Solution Ya-Ting Lin and Chenju Liang* Department of Environmental Engineering, National Chung Hsing University, 250 Kuo-kuang Road, Taichung 402, Taiwan S Supporting Information *

ABSTRACT: Ascorbic acid (AA) mediated electron transfer may induce reductive dechlorination of carbon tetrachloride (CCl4). This study investigated the role of AA in conjunction with the presence of iron minerals over a wide pH range for the reduction of CCl4 in aqueous systems. The results indicate that CCl4 was reduced by AA at a pH of 13 (>pKa2, AA of 11.79) and chloroform (CHCl3) was a transformation byproduct of CCl4. When CCl4 levels were reduced to near complete disappearance, the decrease of CHCl3 was then observed. The degradation rate of CCl4 and also the formation rate of CHCl3 increased with increased AA concentrations. Analysis of reaction kinetics between CCl4 and AA revealed an overall secondorder reaction with a rate constant of 0.253 ± 0.018 M−1 s−1. Furthermore, the reduction rate of CCl4 by AA at pH of 13 could be enhanced with the presence of iron minerals (Fe3O4, Fe2O3, FeOOH, and FeS2). In the absence or presence of iron minerals, the fraction of CCl4 transformed to CHCl3 was less than 1, indicating simultaneous one- and two-electron transfer processes. The end-products of AA at a pH of 13 included threonic acid and oxalic acid. This study highlights the potential of an alkaline AA solution for remediating chlorinated solvents.

1. INTRODUCTION Ascorbic acid (C6H8O6, AA) (also known as Vitamin C) is a water-soluble compound, which is present in most fruits and vegetables. It is also a natural antioxidant and commonly used in agricultural, pharmaceutical, and food industry applications. AA is a dibasic acid, and may serve as a two-electron reductant with a redox potential of −0.06 V.1 AA dissociates in two steps, which are described as first, one hydrogen ion released to form ascorbate anion (C6H7O6−) with a pKa1 value of 4.25 (eq 1), and second, an additional hydrogen ion released to form dianionic ascorbic acid (C6H6O62−) with a pKa2 value of 11.79 (eq 2).1−3 AA and dissociated ions speciation as a function of pH is illustrated in Figure S1 (Supporting Information, SI). C6H8O6 → C6H 7O6− + H+

(1)

C6H 7O6− → C6H6O6 2 − + H+

(2)

studies have researched the application of AA as an electron donor in the reduction of high oxidation state metals and demonstrated that the speed of reduction by AA is pH dependent; preferably at acidic conditions. Through electron transfer process by AA, it has also been demonstrated that the kinetics of methylene blue reduction by AA can be easy and quick under strongly acidic conditions as a “clock reaction”.9,10 However, to date, rather little data exist about the nature and application of AA under basic pH, and also no reports on the reductive capability of AA at elevated pH (e.g., pH above pKa2 value of 11.79) are available in literature. An attempt to discover the reductive capability of AA over a wide pH range for environmental application was one of the objectives in this study. Carbon tetrachloride (CCl4), with a positive carbon oxidation state of +IV corresponding to a species in a highly oxidized form, is considered to be preferentially degraded through reductive pathways (gaining electrons) and hence selected as a model compound for evaluating the reductive capability of AA herein. CCl4, a priority toxic contaminant, has been found in at least 425 of the 1662 most serious hazardous waste sites listed in the United States Environmental Protection Agency’s National Priority List.11 A great number of laboratory

It is well-known that the AA reaction mechanism, and its reducing capability, has a beneficial effect for environmental applications. For example, the point of interest in metal-AA chemistry is the oxidation of AA on the surface of iron oxide where Fe3+ could be reduced to Fe2+.4−6 Moreover, the reductive dissolution rate of Fe3+ is relatively stable in the pH range of 4−6 and decreases markedly above pH 6.4 Also, reduction of hexavalent chromium to trivalent chromium by AA is an environmentally favorable and effective remediation process as the latter species has low toxicity and bioavailability. Evidence has suggested an operative pH be less than 9.7,8 Many © 2013 American Chemical Society

Received: Revised: Accepted: Published: 3299

October 30, 2012 January 12, 2013 February 28, 2013 February 28, 2013 dx.doi.org/10.1021/es304441e | Environ. Sci. Technol. 2013, 47, 3299−3307

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∼45 μm) and hematite (99.5%, ∼45 μm) were purchased from Alfa Aesar. n-Pentane (min. 99.9%) was purchased from Tedia. Acetic acid (100%) was purchased from Taiwan Maxwave Co., Ltd. Methanol (99.9%) was purchased from ECHO Chemical. Water was purified using a Millipore reverse osmosis (RO) purification system. Natural mineral pyrite (97%) was purchased from a local mining company (Ruhnyu, Inc., Tainan, Taiwan) and its characterization was reported in Liang et al.15 Soil used was obtained from farm land located in southern Taiwan and collected from a layer located approximately within 30−100 cm below the ground surface. The characterization of soil was reported in Liang et al.21 The soil was air-dried and sieved (i.e., passed sieve #30 mesh and retained on sieve #200) prior to use. 2.2. Experimental Procedures. Initial experiments focused on investigating the effect of pH on the degradation rate of CCl4 with AA. The pH of RO water was initially adjusted to pH 2, 7, 12, and 13 using NaOH or H2SO4 before the addition of reactants. The CCl4 solution (0.08 mM) was prepared by adding the required amount of pure CCl4 and stirred for 12 h in a 2 L borosilicate reservoir equipped with a Teflon stopper (gastight maintaining no head space inside the reservoir) at 20 °C in a temperature-controlled chamber. Thereafter, a predetermined amount of AA was added to the reservoir and mixed for approximately 2 min reaching a designated AA concentration (10 mM) before filling a series of 60 mL amber screw top PTFE/silicone septum glass reaction bottles (no headspace). All reaction bottles were shaken continuously on an IKA HS 250 reciprocating shaker at 20 °C. Control tests in the absence of AA were also carried out in parallel. Furthermore, two additional sets of experiments were run to investigate (1) the effect of AA concentrations (5, 10, and 30 mM) under a fixed CCl4 concentration of 0.08 mM and (2) the effect of CCl4 concentrations (0.08, 0.32, and 0.56 mM) under a fixed AA concentration of 10 mM. For experiments with the presence of either iron minerals or soils, experimental procedures were similar to those described earlier with the exception that 0.5 g of each iron mineral (Fe3O4, Fe2O3, FeOOH, and FeS2) or 10 g of soil was added into each 60 mL sample bottle prior to filling with CCl4 and/or AA solutions. Three pH levels of 2, 7, and 13 under fixed AA (10 mM) and CCl4 (0.08 mM) concentrations were evaluated. Control tests in the presence of iron minerals or soils without AA were also conducted in parallel. All experiments were conducted in triplicate and averaged data and error range with one standard deviation were presented. 2.3. Chemical Analysis. CCl4 and its associated degradation products in solution were extracted with n-pentane and the extract was analyzed using a gas chromatograph (GC, Agilent 7890 A) with an Agilent DB-624 column (60 m × 0.25 mm i.d.) and a mass spectrometer in electron impact mode (MS, Agilent 5975 C) in accordance with the method for analyzing volatile organic compound in aqueous phase established by the Taiwan National Institute of Environmental Analysis (Method W785.54B). For analysis of AA, aqueous samples were filtered using a polytetrafluorethylene filter (0.2 μm) placed within a stainless syringe holder (Advantec, KS-13) and then measured using a high-performance liquid chromatography (HPLC, Agilent 1100) equipped with UV detector and a ZORBAX Eclipse XDB-C18 (2.1 × 150 mm, 3.5 μm) column. The mobile phase was methanol/0.1% acetic acid (5:95, v/v) at a flow rate of 0.1 mL/min and the effluent was monitored by the UV detector at a wavelength of 265 nm. Transformation

and field studies have employed iron minerals such as zerovalent iron (Fe0), magnetite (Fe3O4), hematite (Fe2O3), goethite (FeOOH), and pyrite (FeS2) in reductive transformation reactions of CCl4 in soil and groundwater systems. The reductive dechlorination of CCl4 by Fe0 occurs by electron transfer coupled with oxidative dissolution of the metal from the Fe0 surface to the adsorbed CCl4. The magnetite, composed of a 1:1 ratio of FeO (Fe2+) to Fe2O3 (Fe3+) could promote CCl4 dechlorination where Fe2+ is the source of electrons.12 Moreover, dechlorination of CCl4 by magnetite is strongly pH dependent over a pH range of 6−10 and the dechlorination reaction proceeded more rapidly as the pH was increased.13 It has been demonstrated that the greater electron density of deprotonated surface sites on magnetite (e.g., ≡FeIIIOFeIIOH0 which is more reactive toward CCl4 than magnetite) at elevated pH may explain the enhanced CCl4 dechlorination rate. Additionally, the electrochemical potential values were found to be more negative with increasing pH, and this was observed to be physically related to the deprotonation of the minerals surface.14 It can be seen that surface bound ferrous species associated with solid minerals, resulting in the formation of reactive surface species, may be more reactive than Fe2+ alone as an electron donor. Moreover, pyrite, a Fe2+-bearing sulfide mineral, can be oxidized with oxygen and Fe3+, which would subsequently release Fe2+ in natural aquatic systems.15 Therefore, in this way, Fe2+ and Fe3+ can be cycled in the presence of both pyrite and Fe3+ (or oxygen) and exhibited the potential to dechlorinate CCl4.16 Hematite and goethite are ferric minerals, which may not be able to directly induce reductive dehalogenation of CCl4. However, sorption or coprecipitation of Fe2+ from aqueous solution onto iron oxides may result in cycling of electrons in the iron mineral-water system, which can be deduced through two sets of mechanisms: (1) the reductive dissolution in the presence of a reducing ligand as electron donor to reduce surface Fe3+ into solution as soluble Fe2+ and (2) the subsequent reduction of Fe3+ by ligand complexed Fe2+ as electron donor where the ligand acts as an electron bridge.5 Hence, many studies have reported successful CCl4 reduction by Fe2+ in the presence of hematite or goethite and dechlorination rates increased with increasing pH from pH 6 to 9.17−19 Rather little is known about how AA acts either alone, or in conjunction with iron minerals, over a wide pH range, in the remediation of organic contaminants in aqueous solution. The industrial and commercial use of AA was mostly applied in formulations of solvents for chemical cleaning and decontamination of oils from metal surfaces.5,20 The major goals of this study are as follows: (i) to investigate the reactivity of AA on reductive transformation of CCl4 under different concentrations and pH conditions (i.e., 2, 7, 12, and 13); (ii) to investigate the effect of AA on CCl4 reduction in the presence of various iron minerals including magnetite, hematite, goethite, and pyrite; and (iii) to evaluate the potential of AA on CCl4 reduction in the presence of natural soils.

2. MATERIALS AND METHODS 2.1. Materials. L-Ascorbic acid (99.7−100.5%), goethite (∼35% Fe), sodium hydroxide (≥99%), and sulfuric acid (95− 97%) were purchased from Sigma-Aldrich. Carbon tetrachloride (min. 99.7%) was purchased from ALPS CHEM Co., Ltd. (Hsinchu, Taiwan). Chloroform (98%) and formic acid (98−100%) were purchased from Merck. Magnetite (97%, 3300

dx.doi.org/10.1021/es304441e | Environ. Sci. Technol. 2013, 47, 3299−3307


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dropped as a result of the H+ released. However, for the pH 13 experiment (pH > pKa2, [OH−] = 0.1 M), a nearly complete dissociation of AA (0.01 M) would generate about 0.02 M [H+] which resulted in a remaining 0.08 M [OH−], corresponding to pH 12.9 and hence the pH remained basic. Furthermore, when comparing ORP data, the redox potential of the system appears decreasing with the increase of pH, and the negative ORP of around −400 mV at pH 13 implies higher reducing power for reductive degradation of CCl4 than those observed under other pH conditions. Collins and Picardal22 studied the degradation rates of CCl4 by dithiothreitol (reducing agent) in the presence of humic acid under a pH range of from 3.6 to 8.7, and reported that the ORP decreased from −90 to −300 mV by increasing the solution pH, which is consistent with accelerated CCl4 degradations. CCl4 was found to be reductively degraded to chloroform (CHCl3) in a solution containing AA at pH 13 (Figure 1). The concentration of CHCl3 was gradually increased, accompanying the decrease of CCl4. Furthermore, when CCl4 decreased to near complete disappearance, a decrease of the CHCl3 concentration was then observed. The rate constants (k2) for the formation of CHCl3 were calculated using the rate law for parallel first order formation of reaction products (eq 4).23 Furthermore, the rate of CHCl3 disappearance may obey eq 5.24 The integrated rate law for CCl4 degradation (eq 3) and the subsequent disappearance of CHCl3 (eq 5) yields eq 6 that characterizes the dynamic transformation of CCl4 and CHCl3.24,25 Table 1 compiles the kinetic data for all experiments. The nonlinear simulations of the data to the model for parallel reactions are also illustrated in Figure 1.

products of AA were qualitatively identified using a Thermo Scientific Surveyor LC plus system and a LTQ linear ion trap mass spectrometer equipped with an electrospray ionization source, a Surveyor MS pump plus and a Surveyor autosampler (LC/MS, Thermo Scientific). Chromatographic separation was done on a ZORBAX Eclipse XDB-C18 column. The mobile phase was methanol/0.01% formic acid (5:95, v/v) at a flow rate of 0.1 mL/min. The specific surface areas of iron minerals were measured using a nitrogen sorption technique at 77 K (BET Sorptometer, CBET-201A, Porous Materials, Inc.). The pH was measured using a pH meter (Thermo Orion 720A+) equipped with a Mettler Toledo Inlab 437 pH combination electrode. Oxidation−reduction potential (ORP) was measured using pH/Conductivity meter (Eutech Instruments CyberScan PC 5000) equipped with a Mettler Toledo Inlab Redox combination electrode.

3. RESULTS AND DISCUSSION 3.1. CCl4 Degradation by AA. 3.1.1. Effect of pHs. In order to assess the effects of pH on the potential AA reductive degradation of CCl4, pH 2, 7, 12, and 13 which cover the range of pKa values were examined. The results show that degradation of CCl4 by AA is strongly pH dependent. At pH 13 AA was capable of reducing CCl4 (see Figure 1) while insignificant

CCl4 degradation: [CCl4]t = [CCl4]0 e−k1t

(3)

CHCl3 formation: k2

CCl4 → CHCl3 + Cl− [CHCl3]t = α[CCl4]0 (1 − e−k1t ); α =

Figure 1. CCl4 transformation to CHCl3 by AA under pH 13. Solid lines show the fits of the data to the model of CCl4 degradation and CHCl3 formation, while dashed lines show the fits of CHCl3 degradation model. Inserted showing rate constant k1 for CCl4, ORP and final pH under different initial pHs. [CCl4] = 0.08 mM, [AA] = 10 mM.

k2 k1

(4)

CHCl3 degradation: d[CHCl3]t = αk1[CCl4]0 − k 3[CHCl3]t dt

CCl4 degradation was observed at other pHs (see Figure S2 for the results obtained at pH 2, 7, and 12, SI). The “inserted figure” in Figure 1 shows a nonlinear relationship between pH and the pseudo-first-order rate constants (k 1 ) of CCl 4 degradations (eq 3), the observed ORP and final pH in solutions. Note that also as shown in Figure 1, the concentration of AA remained nearly constant and therefore the kinetics of the CCl4 degradation fits the pseudo-first-order model (typically R2 > 0.9 and data points within at least 1−3 half-lives used for calculation). The CCl4 degradation rate constant increase with increasing pH (e.g., pH ≈ 13) can be attributed to the formation of dianionic ascorbic acid when the solution pH was above the pKa2 value of 11.79. The results of the control experiments, i.e., CCl4 alone at initial pH 13, exhibited less than 10% CCl4 variation (see Figure 1). It should be noted that, 1 mol of [AA] upon complete dissociation would produce 2 mols of [H+] (see eqs 1 and 2). Hence, it can be seen from the insert in Figure 1 that the solution pH usually

[CHCl3]t

=

(5)

k1[CHCl3]max −k1t (e − e−k3t ) k 3 − k1

[CHCl3]max = α[CCl4]0

(6)

where [CCl4]t and [CHCl3]t are the concentration of CCl4 and CHCl3 versus t, respectively. [CCl4]0 is the initial CCl4 concentration; k1 is the overall pseudofirst-order reaction rate constant for CCl4 degradation, k2 is the rate constant for the formation of CHCl3, and k3 is the rate constant for conversion of CHCl3 to other products; α is the fraction of CCl4 that is transformed to CHCl3, and [CHCl3]max is the maximum production of CHCl3 observed in these systems. 3.1.2. Proposed Reaction Mechanisms. In order to identify possible intermediate products of AA breakdown, the acquired LC/MS spectra are presented in Figures S3−S5 (SI) for the pH conditions of 2, 7, and 13, respectively. Total ion chromatograms in full scan mode within the mass-to-charge ratio (m/z) 3301



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Table 1. CCl4 Transformation Rate Constants at pH 13a AA/CCl4 (mM/mM) 5/0.08 10/0.08 30/0.08 10/0.08 10/0.32 10/0.56 0/0.08 10/0.08 0/0.08 10/0.08 0/0.08 10/0.08 0/0.08 10/0.08

iron mineralb (specific surface area, m2 g−1)

Fe3O4 (20.747) Fe2O3 (24.441) FeOOH (29.671) FeS2 (34.950)

kSA × 10−3 (L d−1 m−2)c

k1 (R2) (d−1)

k2 (d−1)

k3 (d−1)

α (R2)

n.a. 1.661 n.a. 3.222 n.a. 4.171 2.844 3.634

0.124 (0.94) 0.239 (0.96) 0.712 (0.99) 0.239 (0.96) 0.262 (0.99) 0.243 (0.96) n.a. 0.287 (0.99) n.a. 0.656 (0.99) n.a. 1.031 (0.96) 0.828 (0.99) 1.058 (0.96)

0.010 0.028 0.139 0.028 0.039 0.039 n.a. 0.059 n.a. 0.143 n.a. 0.554 0.262 0.686

± 0.022 ± 0.017 ± 0.021 ± 0.017 ± 0.044 ± 0.044 n.a. 0.087 ± 0.068 n.a. 0.171 ± 0.013 n.a. 0.166 ± 0.020 0.173 ± 0.029 0.237 ± 0.011

0.106 (0.99) 0.118 (0.95) 0.195 (0.97) 0.118 (0.95) 0.150 (0.93) 0.162 (0.97) n.a. 0.207 (0.96) n.a. 0.218 (0.99) n.a. 0.537 (0.97) 0.317 (0.95) 0.649 (0.91)

0.091 0.113 0.114 0.113 0.130 0.102

a n.a. = not applicable. bIron mineral (0.5 g/60 mL). ckSA (L d−1 m−2) = (k1 (d−1)/(specific surface area (m2 g−1) × iron mineral mass loading (g L−1)).

Figure 2. (a) Proposed reaction pathways for AA including LC/MS extracted ion chromatograms under pH 13 and (b) CCl4 degradation pathways.

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decarboxylation of THDH may lead to the formation of TA (eq 15).32

range of 85−500 were generated. Thereafter, extractions of parent ion spectra at specific m/z were conducted. Extracted masses of negative ions [M−H]− included possible intermediate products of AA (m/z = 175), at m/z = 173 (dehygroascorbic acid, DHA), m/z = 191 (2,3-diketogulonic acid, DKG), m/z = 207 (4,5,5,6-tetrahydroxy-2,3-diketohexanoic acid, THDH), m/z = 135 (threonic acid, TA), and m/z = 89 (oxalic acid, OA). It is noteworthy that when comparing the relative abundance of m/z in MS spectra among different pH conditions, significant abundances for TA and OA were observed under pH 13 (note that extracted ion chromatograms including compounds with m/z = 175 (AA), m/z = 173 (DHA), m/z = 135 (TA), and m/z = 89 (OA) are presented in Figure 2a). The overall decomposition reactions of AA are initiated in accordance with eqs 1 and 2 based on pHs. Moreover, when alkaline pH was maintained, dissociated AA may further lose one electron via eqs 7 and 8 to induce hydrogenolysis of carbon-chlorine bonds in CCl4 to produce CHCl3; when losing two electrons via eq 9, the alternate CCl4 degradation pathway may undergo dichloroelimination without the formation of CHCl3. The proposed degradation of CCl4 in conjunction with the decomposition of AA is also incorporated in Figure 2b. − C6H6O62 − → C6H6O•− 6 + e

(7)

C6H6O•− 6

(8)

→ C6H6O6 + e

−

C6H6O62 − → C6H6O6 + 2e−

The rate constant (k1) for CCl4 degradation related to AA and dissociated ions may further be expressed as eqs 10 and 11,26,27 k1 = k C6H8O6[C6H8O6 ] + k C6H7O−6 [C6H 7O−6 ] (10)

k1 = {k C6H8O6fC H O + k C6H7O−6 fC H O− + k C6H6O62−fC H O2− } 6

8 6

6

7 6

[AA]tot

6

C6H6O6 + 2OH− → C6H8O8(THDH)

(13)

C6H8O7 + 2CO2 → C4 H8O5(TA) + C2H 2O4 (OA)

(14)

C6H8O8 + 0.5O2 →C4 H8O5 + 2CO2

(15)

6 6

r=−

(11) −

(12)

3.2. Effect of AA and CCl4 Concentrations. Further experiments were conducted to determine the effect of AA and CCl4 concentrations on the degradation of CCl4 specifically at pH 13, (results are shown in Figure S6a,b of the SI), respectively. The kinetic data are summarized in Table 1 and it can be seen that the k1 increased with increased AA concentrations. The higher concentration of AA appears to lead to faster CCl4 degradations and also CHCl3 formations which exhibit a positive correlation that a pathway via hydrogenolysis of CCl4 (formation of CHCl3 as evidence) is dependent on AA concentrations. However, the α value of less than 1 (see Table 1) indicates the occurrence of competing reactions such as reductive dichloroelimination through a dichlorocarbene intermediate (:CCl2, shown in Figure 2b).25 When CHCl3 reached maximum concentrations, the degradation of CHCl3 was observed and the rate constant (k3) for conversion of CHCl3 to other products was found to increase with the increase of AA concentration. It was also seen that the ORP values further decreased with increased concentration of AA. Additionally, three concentration levels of CCl4 were reacted with AA at 10 mM. The results are shown in Figure S6b (SI) and the kinetic data are also presented in Table 1. It can be seen that the k1, k2, and k3 were almost similar in these three reaction systems and this may be because the 10 mM of AA was present in excess to degrade CCl4 and CHCl3. The reduction of CCl4 by AA under pH 13 can be described with the following general rate equation:33

(9)

+ k C6H6O62−[C6H6O62 −]

C6H6O6 + H 2O → C6H8O7 (DKG)

kC6H6O62−

where kC6H8O6, kC6H7O6 , and are the second-order rate constants for the reaction of the CCl4 with AA, ascorbate anion, and dianionic ascorbic ion, respectively; f C6H8O6, f C6H7O6−, and f C6H6O62− are the fraction of concentration as AA, ascorbate anion, and dianionic ascorbic ion, respectively, at specific pH; [AA]tot is the total concentration of AA. On the basis of the k1 data presented in Figure 1 and the fraction of the concentration present as AA, ascorbate anion, or dianionic ascorbic ion (Figure S1 of the SI), the derived constants of kC6H8O6, kC6H7O6−, and kC6H6O62− in accordance with eqs 10 and 11 are 1.62 × 10−5, 1.85 × 10−5, and 2.80 × 10−4 M−1s−1, respectively. Comparison of these values shows that the dianionic ascorbic ion is ∼17 times more reactive than the AA and ascorbate anion. Note that a sample calculation for the derived second-order rate constant and inclusion of these rate constants in AA speciation diagram can be seen in Figure S1 of the SI. Furthermore, DHA can be hydrolyzed causing the opening of DHA-lactam ring to DKG production (eq 12)28 and peroxidation of DHA would also lead to the formation of the THDH (eq 13).29 As a result of carbon 3/carbon 4 cleavage of the DKG structure under alkaline condition,30,31 yields of TA and OA were observed (eq 14). Also, the subsequent

d[CCl4] = k[CCl4]a [AA]b dt

(16)

When the concentration of AA is in excess, eq 16 can be simplied to eqs 17 and 18, where a and b are reaction orders with respect to [CCl4] and [AA]. r = k1[CCl4]a

(17)

k1 = k[AA]b

(18)

where r is a reaction rate, k is a rate constant, and k1 is a pseudofirst-order rate constant as defined in eq 3. By varying the values of [CCl4] and determining reaction rate, the order “a” with respect to [CCl4] can be obtained by a log−log form of eq 17: log r = log k1 + a log[CCl4]

(19)

In a similar way, by varying [AA]0 and determining k1, the order “b” with respect to [AA]0 can also be determined by a log−log form of eq 18: log k1 = log k + b log[AA]0

(20)

To avoid complications from subsequent reactions with CCl4 degradation byproducts, an initial rate method was used here and eq 19 can be expressed as follows: 3303


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constant, i.e., log k = 1.339 ± 0.030, and k = 0.253 ± 0.018 (M−1 s−1). Hence, eq 16 can be rewritten as eq 22.

(21)

The initial reaction rates (r0) were estimated as the tangent to the CCl4 concentration−time curve. As shown in Figure 3,

r = 0.253 ± 0.018M−1s−1[CCl4]1.014 ± 0.018 [AA]0.978 ± 0.015 (22)

3.3. CCl4 Degradation by AA in the Presence of Different Iron Minerals. Iron(III) oxides may be dissolved mainly via three mechanisms: dissolution by acid, detachment of Fe3+ by specifically adsorbed ligands, and reduction by reductant.5 Elucidation of the effect of AA and pHs (i.e., pH 2, 7, and 13) on different iron minerals (Fe3O4, Fe2O3, FeOOH, and FeS2) for degradation of CCl4 were conducted. It is conspicuous that only pH 13 revealed significant CCl4 degradations by AA (see Figure S7of the SI). Also, the much lower ORP, as evidence of reducing conditions, was observed when there was an increased solution pH. Even though the dissociation of AA may induce reductive, chelating, and acidic dissolution of iron oxides to dissolved Fe2+, which also acts as a reductant,5 the reduction of chlorinated methane such as CCl4 seems slow. However, Pecher et al.19 reported that reduction rates of polyhalogenated methanes could be enhanced with the assistance of surface-bound ferrous ion and surface precipitation to iron oxides when the pH was increased from 6 to 8.9. Thus, it was speculated that iron minerals could be reduced by AA and Fe2+ could be adsorbed on the surface of iron minerals at elevated pH (e.g., pH 13) to accelerate CCl4 degradation. Also, when compared to the results obtained in the absence of AA, iron minerals under different pHs were unable to significantly induce reduction of CCl4, except for the observed degradation of pyrite at pH 13. Pyrite consists of two surface sites (>FeSS and >SSFe), the pHzpc of pyrite was around 2.7.34 When pH > pHzpc, the surface charge of pyrite is negative and >SSFeO− and >FeSS− surface sites are present. The >FeSS− sites dominate at neutral or alkaline pHs and are more hydrophobic than >SSFe− site. Thus, the hydrophobic molecules such as CCl4 could be

Figure 3. Plot of initial rates versus initial concentration under pH 13 for three kinetic experiments ([CCl4] = 0.08, 0.32, and 0.56 mM), [AA] = 10 mM); and plot of k1 versus initial concentration of AA under pH 13 ([CCl4] = 0.08 mM), [AA] = 5, 10, and 30 mM).

the slope a = 1.014 ± 0.018 (R2 = 0.998) was calculated (based on eq 21) through a linear regression of the log of initial rates versus the log of initial CCl4 concentration. A slope of b = 0.978 ± 0.015 (R2 = 0.999) was determined from a plot of the log of k1 versus the log of [AA]0 (Figure 3). The results from Figure 3 demonstrate that the reaction between CCl4 and AA is nearly a second-order reaction with a = 1.014 ± 0.018 and b = 0.978 ± 0.015. The y-intercept of the trend line (1.339 ± 0.030) is the log of the second-order rate

Figure 4. CCl4 transformation to CHCl3 by AA as a function of reaction time in the presence of (a) magnetite, (b) hematite, (c) goethite, and (d) pyrite under pH 13. Solid lines show the fits of the data to the model of CCl4 degradation and CHCl3 formation, while dashed line shows the fits of CHCl3 degradation model. [CCl4] = 0.08 mM, [AA] = 10 mM, iron mineral = 0.5 g/60 mL. 3304



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making CHCl3 as an intermediate should be a focus of future studies. 3.4. CCl4 Degradation in Soil Slurries. The potential application of AA solutions for remediating CCl4 contamination was preliminarily evaluated in a soil slurry system. As shown in Figure 5, CCl4 can be degraded by AA in the soil

preferentially adsorbed on sulfide sites for subsequent electron transfer.35−38 FeS2 + 11H 2O → Fe(OH)3 + 2SO24 − + 19H+ + 15e− (23) 2+

In alkaline AA solution, when Fe is released, e.g., being reduced or chelated by the presence of AA, ferrous ions tend to precipitate as ferric hydroxide on the pyrite surface and produce electrons in accordance with eq 23.15 Several studies indicated that the degradation of CCl4 by surface-bound Fe2+ on hematite,19 goethite,17,18 and magnetite13,19,39 and zerovalent iron40−42 could also take place in a similar manner to pyrite. Furthermore, assessment of the transformation of CCl4 to CHCl3 by AA in the presence of different iron minerals under pH 13 as a function of reaction time was conducted and Figure 4 shows the results of CCl4 transformation. The determined rate constants are also presented in Table 1. To fairly compare the effects of surface area of different iron minerals on rate constants, surface area normalized rate constants (kSA)43 were calculated and reported in Table 1. Note that the specific surface areas for Fe3O4, Fe2O3, FeOOH, and FeS2 were determined to be 20.747, 24.441, 29.671, and 34.950 m2/g, respectively. An order of effect of iron minerals on the kSA of CCl4 degradation was FeOOH > FeS2> Fe2O3 > Fe3O4. The trend in kSA values is generally comparative to their corresponding k1 values. FeOOH and FeS2 exhibit the most reactivity in transforming CCl4 in the alkaline AA reaction system. It can be seen that the rate constants (i.e., k1) of CCl4 degradation by AA in the presence of different iron minerals and the fractions of CHCl3 formed (i.e., α) were generally higher than the rate determined in the absence of iron minerals. Also, the formation rate of CHCl3 (i.e., k2) would be increased when the degradation rate of CCl4 was increased. Among these iron minerals tested, the presence of goethite or pyrite (see Figure 4c,d) tends to promote more one-electron and less twoelectron transfers than the presence of magnetite and hematite, as evidenced by higher CHCl3 generation. These results are similar to a study that investigated the transformation of CCl4 to CHCl3 by goethite associated with Fe2+ at pH 7.2, and reported that both one- and two-electron processes took place in parallel.19 It should be noted that the presence of AA can interact with pyrite, and thus the degradation rate of CCl4 by AA in the presence of pyrite was higher than the rate obtained in the absence of pyrite, as explained earlier. Chloroform is the major byproduct observed during the course of CCl 4 degradation by alkaline AA. However, the α value, as one measure of the branching between the one- and two-electron pathways for reduction of CCl4,24,42 varies with different alkaline AA reaction systems. For example, alkaline AA alone usually resulted in α values less than 0.2, while alkaline AA in the presence of iron minerals resulted in α values ranging from 0.2 to 0.6. A lower α value indicates a favorable two-electron reaction pathway over the one-electron pathway and less undesirable CHCl3 generation.42 Nevertheless, CHCl3 would still be degraded in the alkaline AA system with or without the presence of iron minerals, as seen in the kinetic experiments. Furthermore, byproducts other than CHCl3 are still unclear and further studies to determine all of the byproducts under a variety of experimental conditions are required. Emphasis on optimizing the reduction of CCl4 to nonchlorinated products such as formate, carbon monoxide, and methane39,42 via the two-electron transfer, or further carbene reduction (i.e., dichlorocarbene (:CCl2)) pathways (see Figure 2b) without

Figure 5. CCl4 degradation by AA in the presence of soil under different pHs. [CCl4] = 0.08 mM, [AA] = 10 mM, soil = 0.5 g/60 mL. Reaction time = 7 d.

slurry system at pH 13. The observed remaining CCl4 (11%) in the soil slurry system after the reaction is analogous to 20% observed in the aqueous system (see Figure 1). Note that in control experiments in the absence of AA, it was seen that overall extraction recoveries of CCl4 ranged from 78% to 85% and the loss of CCl4 may have occurred due to possible biodegradation activities or bonding to the soil matrix. The native soil minerals such as iron oxides might be reduced by alkaline AA to form surface-bound Fe2+ on soil minerals for reduction of CCl4. Also, the soil organics such as humic acid may act as electron transfer mediators. 44 Moreover, CHCl3 was detected only under pH 13 in the presence of AA, as evidence of CCl4 degradation, while no CHCl3 was detected under all other experimental conditions. On the basis of these results, it is concluded that alkaline AA solution presents a significant potential for remediating chlorinated solvent contamination. Moreover, pH adjustment for the alkaline AA solution after treatment may be needed to meet specific environmental conditions, e.g., neutralization of the alkaline solution using more AA or other organic/inorganic acids.

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ASSOCIATED CONTENT

S Supporting Information *

Correlation between the second-order rate constants for AA and dissociated ions and AA speciation under different pH values (Figure S1); CCl4 degradation by AA under different pHs (Figure S2); LC/MS spectra of AA and its byproducts under pH 2 (Figure S3); LC/MS spectra of AA and its byproducts under pH 7 (Figure S4); LC/MS spectra of AA and its byproducts under pH 13 (Figure S5); influence of AA and CCl4 concentration transformation under pH 13 (Figure S6); and CCl4 degradations by AA in the presence of different iron minerals under pHs (Figure S7); and additional calculations and references. This material is available free of charge via the Internet at http://pubs.acs.org. 3305



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AUTHOR INFORMATION

Corresponding Author

*Tel.: +886-4-22856610; fax: +886-4-22856610; e-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS

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REFERENCES

This study was funded by the National Science Council of Taiwan under Project No. NSC 100-2221-E-005-006-MY3. The authors acknowledge Prof. Lee, Maw-Rong, and Dr. Lee, Ren-Jye from the Department of Chemistry at National Chung Hsing University for LC/MS analysis.

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Carbon Tetrachloride Degradation by Alkaline Ascorbic Acid Solution

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