Thermally activated peroxydisulfate in the presence of additives: A clean method for the degradation of pollutants

Chemosphere 75 (2009) 1405–1409

Contents lists available at ScienceDirect

Chemosphere journal homepage: www.elsevier.com/locate/chemosphere

Technical Note

Thermally activated peroxydisulfate in the presence of additives: A clean method for the degradation of pollutants Verónica C. Mora a, Janina A. Rosso a, Galo Carrillo Le Roux b, Daniel O. Mártire a, Mónica C. Gonzalez a,* a

Instituto de Investigaciones Fisicoquímicas Teóricas y Aplicadas (INIFTA), Facultad de Ciencias Exactas, Universidad Nacional de La Plata, CONICET. C.C. 16, Suc. 4, (1900) La Plata, Argentina b Departamento de Engenharia Quimica, Escola Politécnica da Universidade de São Paulo, Av. Prof. Luciano Gualberto, Trav. 3, 380, P.O. Box 05508-900 São Paulo, SP, Brazil

a r t i c l e

i n f o

Article history: Received 24 November 2008 Received in revised form 11 February 2009 Accepted 12 February 2009 Available online 13 March 2009 Keywords: Carbon dioxide radical anions Sulfate radicals Trichloroacetic acid Waste water treatment Activated peroxydisulfate

a b s t r a c t The kinetics and mechanism of the thermal activation of peroxydisulfate, in the temperature range from 60 to 80 °C, was investigated in the presence and absence of sodium formate as an additive to turn the oxidizing capacity of the reaction mixture into a reductive one. Trichloroacetic acid, TCA, whose degradation by a reductive mechanism is well reported in the literature, was used as a probe. The chemistry of thermally activated peroxydisulfate is described by a reaction scheme involving free radical generation. The proposed mechanism is evaluated by a computer simulation of the concentration profiles obtained under different experimental conditions. In the presence of formate, SO 4 radicals yield CO 2 , which are the main species available for degrading TCA. Under the latter conditions, TCA is more efficiently depleted than in the absence of formate, but otherwise identical conditions of temperature and ½S2 O2 8 . We therefore conclude that activated peroxydisulfate in the presence of formate as an additive is a convenient method for the mineralization of substrates that are refractory to oxidation, such as perchlorinated hydrocarbons and TCA. This method has the advantage that leaves no toxic residues. Ó 2009 Elsevier Ltd. All rights reserved.

1. Introduction Activated sodium peroxydisulfate has the potential to in situ destruct many organic contaminants commonly encountered in soil and groundwater (House, 1962; Liang et al., 2003; Huang et al., 2005). These properties summed to its safe handling and its solubility in water makes it an excellent additive for waste treatment. Although a strong oxidant, peroxydisulfate anion, S2 O2 8 , slowly reacts with many organics (Osgerby, 2006). It can be chemically (i.e.: via transition metals (Anipsitakis and Dionysiou, 2004)), photochemically or thermally activated to generate the stronger oxi 2  dant sulfate radicals, SO 4 with a redox potential E ðSO4 =SO4 Þ ¼ 2:6 V (Wardman, 1989). Generation of these radicals can significantly accelerate the kinetics of oxidation in a wide range of matrix conditions. Thermally activated peroxydisulfate has been demonstrated to be able to decompose many contaminants in aqueous systems, including chlorinated methanes, ethanes and ethenes (Killian and Bruell, 2003; Liang et al., 2003, 2004; Huang et al., 2005; Waldemer et al., 2007). Substrate degradation rates increased with increasing temperature and oxidant concentration and were strongly affected by background matrix components, such as carbonate and pH * Corresponding author. Tel.: +54 221 425 74 30; fax: +54 221 425 46 42. E-mail addresses: [email protected] (G. Carrillo Le Roux), [email protected]. edu.ar (M.C. Gonzalez). 0045-6535/$ - see front matter Ó 2009 Elsevier Ltd. All rights reserved. doi:10.1016/j.chemosphere.2009.02.038

(Liang et al., 2006, 2007). The refractory to oxidation CCl4 substrate was not degraded by activated peroxydisulfate when treated individually in deionized water, but was observed to be degraded if treated in a mixture of contaminants (Huang et al., 2005). We suspect that CCl4 depletion is possible under the latter conditions due to the occurrence of a reducing mechanism. Despite there is consensus on the complexity of the reaction mechanisms involved (Liang and Bruell, 2008), the kinetic information necessary for the optimization and efficiency evaluation of in situ application of activated peroxydisulfate is not abundant in the literature. Oxidation of formic acid or its conjugated base by sulfate radicals (Ross et al., 1998) leads to the formation of carbon dioxide rad ical anions, CO 2 , with reduction potential EðCO2 =CO2 Þ ¼ 2:0 V vs. SHE (Wardman, 1989). Therefore, addition of formic acid may turn the oxidative environment established by activated peroxydisulfate in a reductive one. Trichloroacetic acid (TCA) was thought as an interesting sensing molecule to investigate the reductive capacity of the reaction mixture since it is refractory to mineralization by oxidative processes at room temperature and was shown to quantitatively release organic chlorine as chloride ions mainly by reaction with reducing species (David Gara et al., 2007). In the present study we report a kinetic and mechanistic investigation on the thermal activation of peroxydisulfate in the presence and absence of potassium formate as an additive. The aim of these investigations is to obtain useful information necessary

V.C. Mora et al. / Chemosphere 75 (2009) 1405–1409

to evaluate the use of activated peroxydisulfate for the mineralization of substrates that are refractory to oxidation.

0.5

0.0

3.1. Experiments with trichloroacetic acid and peroxydisulfate Experiments performed at 80 °C with air-saturated solutions of pH in the range from 3 to 4 containing 1–5 mM TCA and 2.5– 25 mM Na2S2O8 showed considerable TCA degradation to carbon dioxide and chloride ions in less than 2 h. All TCA, TOC and peroxydisulfate time profiles could be well fitted to a first order law with apparent rate constants kapp(TCA), kapp(TOC), and kapp ðS2 O2 8 Þ, respectively. Chloride anion formation was fitted to an exponential raise of the form a(1  exp(kt)), with k = kapp(Cl) and ‘‘a” the final chloride anion concentration. Experiments performed under similar experimental conditions but in the absence of molecular oxygen show negligible amounts of Cl formation. Fig. 1a shows the time profiles of [TCA]/[TCA]0, TOC/TOC0, and [Cl]/(3[TCA]0) for one typical experiment (the subscripts ‘‘0” stand for initial values and (3[TCA]0) is the maximum chloride concentration that can be formed). From the fitting of the traces, the apparent rate constants kapp(TCA) = kapp(TOC) = kapp(Cl) = (8 ± 1)  104 s1 were obtained. The chlorine atom mass balance ([Cl] + (3[TCA]))/(3[TCA]0) ffi 1 (also shown in Fig. 1a) and the

-2.1 -2.4

0

200

a 0

15

30

45

60

3.0

2.0

1.0

b 0

50

[Cl ] / [TCA] or [TCA] / [TCA] 0 0

2.1. Analytical methods

3. Results and discussion

-1.8

time / min

0.0

100

1.0 100

6

After mixing the reactants, 5 mL sample aliquots were periodically collected from the reactor, cooled in an ice-bath and thermostatized to 25 °C before HPLC, TOC and chloride anion analysis. Chloride ion was analyzed by a selective chloride electrode (Phoenix Electrode Company). Detection limits were [Cl] ffi 0.01 mM, see Supporting Information (SI) Note 1. TCA and other Cl containing acids were analyzed by HPLC with a Hewlett-Packard HPLC model 1050 (Ti series) chromatograph with multiwavelength detection, a C18 Restek Pinacle II column (particle size 5 lm, 2.1 mm i.d.  250 mm) and using a 25/75/ 0.1% (v/v) methanol/water/H3PO4 mixture as eluent at 0.5 mL min1 constant flux. Detection limits were 0.2 mM for TCA. The same method allows determination of peroxydisulfate with a detection limit of 1 mM. TOC was measured with a Shimadzu TOC-5000A analyzer. Detection limits were 1 ppm C.

log [Na2S2O8]

Relative concentration

TCA (Fluka, puriss), KC8H5O4 (Potassium biphtalate, UCB), KCHO2 (Carlo Erba p.a.), Na2S2O8 (Riedel-de Haen, AG), NaNO3 and Na2CO3 (Merck) and KCl, K2HPO4 and KH2PO4, (all J.T. Baker ACS), were used without further purification. Distilled water (>18 MX cm, <20 ppb of organic carbon) was obtained from a Millipore system. The temperature was controlled to ±1 °C with a Grant model GD 1200 thermostat. The pH of the solutions used in experiments with Na2S2O8 and TCA was at their original pH in the range 3–4 and therefore TCA (pKa = 0.77) was present as its conjugate base. The pH of the solutions with Na2S2O8, TCA and KCHO2 was controlled to 6–7 by addition of 0.4 M phosphate buffer such that mainly the conjugated base of formic acid is present (pKa = 3.75 at room temperature). The ionic strength, of the solutions was 0.9–1 M. Freshly prepared solutions of the reactants were separately heated for 15 min to the desired temperature and mixed in a 250 mL volume reactor immersed in the thermostat bath. The difference in the peroxidisulfate concentration before and after heating was of order of, or lower than, the experimental error in the determination of its concentration by HPLC.

[TCA] or [Cl ] / mM

2. Materials and methods

1.0

kapp x 10 /Ms-1

1406

0.5

50 0 0

[S2O82-] / mM

10

c

0.0 0

15

30

45

60

time / min Fig. 1. (a) [TCA]/[TCA]0 (r), TOC/TOC0 (d), [Cl]/(3[TCA]0) (j), and chlorine atom mass balance ([Cl] + (3[TCA]))/(3[TCA]0) (}) in experiments at 80 °C with solutions of pH 3–4 containing 1 mM TCA and 10 mM Na2S2O8. Inset: semi logarithmic plot of [Na2S2O8] vs. time for experiments with 1 mM TCA and different initial Na2S2O8 concentrations: 2.5 mM (s), 7.5 mM (}), 10 mM (h), and 20 mM (4). (j) and (N) stand for experiments at 80 °C with 2 mM initial TCA and [Na2S2O8]0 = 10 mM and 20 mM, respectively. Dotted lines stand for a fit to a first order rate law. Solid curves represent computer simulations (refer to text). (b) TCA depletion (open symbols) and chloride ion evolution (full symbols) in experiments at 80 °C with 10 mM Na2S2O8 buffered solutions of pH 6–7 containing (s) 1 mM TCA and 1 mM KHCO2, (4) 1 mM TCA and 10 mM KHCO2, and (h) 2.3 mM TCA and 10 mM KHCO2. (c) [TCA]/[TCA]0 (full symbols) and [Cl]/[TCA]0 (open symbols) time profiles obtained from experiments at 80 °C with buffered solutions of pH 6–7 containing 1 mM TCA, 10 mM KCHO2, and either 2.5 (s) or 10 mM Na2S2O8 (D). Inset: plot of kapp vs. [Na2S2O8]. Error bars in the kapp values are on the order of the symbol size, unless otherwise specified.

identical rates of TOC/TOC0 and TCA/TCA0 depletion indicate that formation of chlorine-containing organic products may be at trace quantities. TCA and TOC depletion, and chloride anion rise rate increase with increasing [Na2S2O8]0 (see SI Fig. 1). Depletion of peroxydisulfate is significant during the experiment time, and therefore [Na2S2O8] cannot be taken as a constant value. Fig. 1a inset shows a semilogarithmic plot of [Na2S2O8] vs. time for experiments with different initial concentrations of TCA and S2 O2 8 . The straight lines

1407

V.C. Mora et al. / Chemosphere 75 (2009) 1405–1409 4 1 with similar slopes (average kapp ðS2 O2 s Þ 8 Þ ¼ ð1:0  0:5Þ  10 indicate that peroxydisulfate decay follows a first order rate law, in line with literature reports on peroxydisulfate thermal activation (Johnson et al., 2008). The effect of temperature on the rate of TCA and TOC depletion, and chloride anion formation was studied in experiments with solutions of pH 3–4 containing 1 mM TCA and 2.5 mM Na2S2O8 performed at 25, 60, 70 and 80 °C. Experiments at 25 °C showed no TCA degradation after 7 d. Within the experimental error, kapp(TCA) = kapp(TOC) = kapp(Cl) holds for T P 60 °C (experiments not shown). The slope of the Arrhenius plots for kapp(TCA), kapp (TOC), and kapp(Cl) yield Ea = (115 ± 10) kJ mol1, see SI Fig. 2. To determine the participation of sulfate radicals formed after peroxydisulfate activation, Reaction (1) in Table 1, the effect of rad 2  ical scavengers such as HCO 2 =H2 CO2 ; HCO3 , and HPO4 =H2 PO4 on the depletion rate of TCA was determined (see SI Fig. 3). TCA deple2  tion rates decrease in the presence of HCO 3 and HPO4 =H2 PO4 , as expected from their reaction with sulfate radicals leading to the formation of the weaker oxidants carbonate and phosphate radicals (Ross et al., 1998; Mártire and Gonzalez, 2001). On the other hand, reaction of sulfate radicals with HCO 2 =H2 CO2 leads to an increase in the decay rate of TCA (vide infra). Therefore, a radical mechanism takes place and any elemental thermal reaction between TCA and peroxydisulfate anions should be of lesser significance. Sulfate radicals react with chlorinated aliphatic carboxylic acids and their conjugated base by H abstraction and electron transfer, respectively, to yield carboxyl radicals which eliminate CO2 (Madhavan et al., 1978). Reaction (2) in Table 1 assumes a similar reaction path for TCA. Since chloride anions are released immediately after TCA depletion and since chloro-containing organic by-products are formed in  trace quantities (if any), the trichloromethyl radicals, CCl3 , generated in Reaction (2) should further undergo thermal reactions lead ing to their mineralization. CCl3 radicals are rapidly oxidized by dissolved molecular oxygen (7.3  102 mM in air-saturated solutions at 80 °C) with no stable chloro-containing intermediates produced (Kin and Choi, 2002), Reaction (3). The latter suggestion is further supported by the experiments performed in the absence of molecular oxygen showing no Cl anion formation. Any reaction between peroxydisulfate and the organic radicals should be of little significance since peroxydisulfate decomposition does not depend, within a 30% error, on TCA concentration.

Recombination of sulfate radicals, Reaction (4), their reaction  anions (Huie and Clifton, 1990; Herwith water, S2 O2 8 , and Cl mann et al., 1995; Alegre et al., 2000; Yu et al., 2004), Reactions (5), (6) and (7), respectively, diminish the efficiency of TCA oxidation by SO 4 . Within the experimental error, depletion of chloride anion by sulfate radicals is not observed until TCA is quantitatively depleted (see SI Fig. 4). However, the occurrence of the reversible reaction of SO 4 with chloride anions (Hermann et al., 1995), Reactions (7) and (7), cannot be neglected. Chlorine atoms formed after Reaction  (7) reversibly react with Cl anions to yield Cl2 radical anions (Yu et al., 2004). The chemistry involved in aqueous solutions with [Cl] 6 3 mM, as in our reaction system, is mainly that of Cl (Alegre et al., 2000). Reaction of Cl with water involves a series of reversible processes leading to the final formation of Cl and HO. When HO radicals are completely removed, no equilibrium is established and the reactions of Cl with water follow simple first order kinetics (Alegre et al., 2000), as shown in Reaction (8). Possible reactions removing HO radicals in the reaction system are ( Reaction (9)). Reaction recombination and reaction with S2 O2 8 of chlorine atoms with peroxydisulfate, Reaction (10), was also reported to deplete Cl (Yu et al., 2004). The complex reaction mechanism may be validated by a computer simulation (see SI Note 2) of the concentration profiles of TCA, Cl, and peroxydisulfate, of a set of eight experiments analyzed globally. Input parameters are the initial concentration of reactants. The program incorporates the set of reactions shown in Table 1, which includes a minimum number of reactions. A discussion of the values employed for some reported rate constants may be found in SI Note 3. The good fit of the global analysis is shown by the solid lines in Fig. 1a for a typical experiment (SI Fig. 5 shows the fitting of some other experiments). The set of reactions shown in Table 1 is those proved to be of significance to the overall reaction (see also SI Note 4). Table 1 also depicts the best set of values obtained for the rate constants. Taking k2 < 2  106 M1 s1 at room temperature (Ross et al., 1998) and k2 (0.4–1)  108 M1 s1 at 80 °C, an activation energy of 43–64 kJ mol1 is estimated for Reaction (2), on the order reported for chlorinated ethanes (Liang et al., 2004; Huang et al., 2005). Reaction (2) involves an electron transfer from the substrate to the radical, however, the pre-exponential factor of 1014–1017 observed may indicate that the reaction involves short-lived adducts

Table 1 Manifold of reactions taking place in the thermal activation of peroxydisulfate in the presence of TCA and the corresponding reaction rate constants k at 80 °C. Activation energies, Ea, calculated from the results herein presented and those reported in the literature are also shown. Ea (kJ mol1)

k at 80 °C  S2 O2 8 ! 2SO4  2  þ CCl -CO SO 3 4 2 ! SO4 þ CCl3 -CO2 CCl3  CO2 !  CCl3 þ CO2 ðþ2H2 O1=2O2 Þ  CCl3 þ O2 !  OOCCl3 !! CO2 þ 3HCl  2 2SO4 ! S2 O8  2 þ SO 4 þ H2 O ! HO þ SO4 þ H 2  2 SO þ S O ! S O þ SO 2 8 2 8 4 4   2 SO 4 þ Cl ! SO4 þ Cl   2  SO4 þ Cl ! SO4 þ Cl   Cl þ H2 O ! Hþ þ Cl þ HO   HO þ S2 O2 8 ! S2 O8 þ HO   ! Cl þ S2 O Cl þ S2 O2 8 8 a b c d e f g h

(0.7–1)  10 s (0.4–1.7)  108 M1 s1b Fast

96 ± 40a 43–64b

(1) (2)

1  1010 M1 s1c 5  108 M1 s1c k5  [H2O] = 1300 s1c (0.7–1)  107 M1 s1b 2.5  108 M1 s1e P2.5  108 M1 s1e,h 2  105 s1g (3–5)  108 M1 s1b 9  106 M1 s1g

<15b <5d 9.2e 35f 0g

(3) (4) (5) (6) (7) (7) (8) (9) (10)

4

þ 1=2H2 O2

Retrieved by the simulations and in agreement with k1 and Ea1 measured in this work. Obtained from the simulation of the experiments. Taken from Ross et al. (1998) and corrected for the temperature, see text. David Gara et al. (2008). Hermann et al. (1995). This work, in agreement with Ea6 + Ea7 = 49 kJ mol1 (David Gara et al., 2008). Yu et al. (2004). The simulation is not sensitive to this reaction if k2 > 4  107 M1 s1.

1a

40b

1408

V.C. Mora et al. / Chemosphere 75 (2009) 1405–1409

produced in an addition mechanism followed by subsequent transformation. A similar behavior was reported for several SO 4 reactions with inorganic anions (Hermann et al., 1995). A plot of the concentration of total reacted peroxydisulfate at a given time vs. the concentration of TCA degraded, yields information on the system efficiency in depleting TCA (see SI Fig. 6). A system. slope of 0.6 is observed for the 1 mM TCA/10 mM S2 O2 8 Experiments with added 1 mM KHCO2 show slopes >0.6, indicating a more efficient depletion of TCA per mol of oxidant consumed in the presence of formate. Experiments with added 10 mM Na2CO3 show slopes <0.6, as expected from the scavenging of SO 4 by HCO 3 . Therefore, TCA oxidation is limited by Reaction (2), as also reported for the degradation of chlorinated ethenes (Waldemer et al., 2007). The activation energy of 115 kJ mol1 found for kapp(TCA), vide supra, is due to the combination of the activations energies of Reactions (1) and (2), and those reactions competing for sulphate radicals (mainly Reactions (5)–(7)). According to the mechanism in Table 1, kapp ðS2 O2 8 Þ ffi k1 þ  , with ss the steady state concentration. Taking k6  ½SO 4 ss 9 mM for a typical experiment such as that ½SO 4 ss ð1—1:5Þ  10 shown in Fig. 1a, see SI Note 4, then Reaction (6) contributes 20% to peroxydisulfate depletion. This contribution is on the same order of magnitude than our experimental error in the determination 2 of kapp ðS2 O2 8 Þ, and therefore, differences in kapp ðS2 O8 Þ for the different experiments could not be appreciated. 3.2. Experiments with TCA and formic acid Fig. 1b and c shows the concentration profiles for TCA depletion and chloride anion formation for experiments with different initial concentrations of TCA, peroxydisulfate, and KHCO2 in buffered solutions of pH 6–7. For each experimental condition used, TCA depletion and chloride ion formation rates are, within the experimental error, coincident up to reaction times of circa 20 min, therefore indicating that only one organic chlorine atom is eliminated as chloride ion per TCA consumed. Further thermal reactions may quantitatively convert the remaining organic chlorine to chloride ions as observed for reaction times higher than 20 min. Quantitative recovery of chloride anions strongly depends on the concentrations of formate (see Fig. 1b). TCA decay and initial Cl formation rates follow zero order rate law kinetics in TCA, as shown by the straight lines in Fig. 1b and c. The slopes of these lines yield an apparent rate constant, kapp, which was observed to be independent of TCA and KHCO2 (see Fig. 1b) in the concentration range from 0.05 to 3 mM and from 1 to 100 mM, respectively. However, kapp linearly depends on peroxydisulfate concentration as shown in Fig. 1c inset. From the slope of the plot of kapp vs. [Na2S2O8], an absolute rate constant kTCA = (1.2 ± 0.2)  104 s1 is obtained at 80 °C. Eq. (1) shows the experimental rate law obtained for TCA decay



d½TCA ¼ kTCA  ½Na2 S2 O8  dt

ð1Þ

TOC depletion also follows a pseudo zero order rate law strongly depending on peroxydisulfate concentration. Experiments performed in the temperature range from 65 to 87 °C with solutions containing 1 mM TCA, 10 mM formate and variable amounts of peroxydisulfate show, at a given temperature, a linear dependence of kapp with ½S2 O2 8  (vide supra) only for experiments with T P 75 °C. Therefore, Eq. (1) applies only in this temperature range. The calculated values of kTCA are (5.0 ± 1.5)  105 s1, (1.2 ± 0.2)  104 s1 and (1.5 ± 0.1)  104 for 75, 80 and 87 °C, respectively. An Arrhenius plot for kTCA yields an activation energy of 96 ± 40 kJ mol1. In the presence of formate, SO 4 produced in Reaction (1) are scavenged to yield CO 2 , Reaction (11) in Table 2. The latter rad-

Table 2 Manifold of reactions taking place in the thermal activation of peroxydisulfate in the presence of formate and TCA.  2  þ SO 4 þ HCOO ! SO4 þ CO2 þ H    CO 2 þ CCl3  CO2 ! CO2 þ Cl þ CCl2 -CO2 k12 ð80  CÞ 1:5  109 M1 s1   2 CCl2 CO 2 ! O2 CCCl2 CCl2 CO2    CCl2  CO þ CO ðþH OÞ !! 2Cl þ CO2 þ Hþ þ COHCO 2 2 2 2     CCl2  CO þ HCOO ! HCCl  CO 2 2 2 þ CO2  CO þ O ! O þ CO 2 2 2 2 k16 ð80  CÞ 2  1010 M1 s1

(11) (12) (13) (14) (15) (16)

icals may initiate the reductive dehalogenation of TCA, Reaction (12), which involves the release of one organic chlorine atom as chloride anion, in agreement with the simultaneous release of one chloride anion per molecule of TCA depleted. Further thermal reactions of  CCl2 CO 2 radicals formed after Reaction (12) release the remaining organic chlorine as Cl, as observed at longer reaction times. The  CCl2 CO 2 radicals may recombine to yield tetrachlorobutanedioic acid, be further reduced by CO 2 to as final products yield CO2, HCl and small quantities of C2 O2 4 (David Gara et al., 2007), or react with formate to yield dichloroacetic acid, Reactions (13)–(15), respectively. Reaction (15) may be the reason for the observed incomplete conversion of organic chlorine to chloride anions in experiments with high concentrations of formate (see Fig. 1b). Reaction of molecular oxygen with CO 2 radicals, Reaction (16), is very fast and may therefore compete with TCA for CO 2 radicals, Reaction (12). However, our experiments under argon- and airsaturation showed no significant differences and Reaction (16) may be neglected under our experimental conditions. Taking [O2] = 7.3  102 mM in air-saturated solutions at 80 °C, and k16 6 2  1010 M1 s1, the upper limit restriction for a diffusion controlled reaction at 80 °C (see SI Note 3), the condition k12  [TCA] P 1.5  106 M1 s1 applies, and therefore k12 (80 °C) P 1.5  109 M1 s1. The solution of the differential algebraic equation system of the reaction mechanism formed by Reactions (1), (11), and (12) assuming the steady state condition for sulfate and carbon dioxide radical anions, yields an expression for the decay rate of TCA identical to Eq. (1) with kTCA = k1. In fact, k1 = (1.2 ± 0.2)  104 s1 is in agreement with the value retrieved by the computer simulations of the experiments in the absence of formic acid (see Table 1) and Ea1 = 96 ± 40 kJ mol1 agrees with the reported energy for the homolysis of peroxydisulfate (Brusa et al., 2000). Therefore, TCA degradation rate by activated peroxydisulfate in the presence of formic acid is limited by the decomposition of peroxydisulfate. 3.3. General conclusions obtained from the kinetic analysis Some general conclusions of importance for future activated peroxydisulfate applications may be obtained from the kinetic analysis. The depletion efficiency for a given contaminant, S, when treated individually with activated peroxydisulfate in aqueous solution is given by the ratio of rate constants of the substrate reaction with   SO 4 ; kðS þ SO4 Þ, and those of competing reactions consuming SO4 (mainly reactions (5)-(7) for solutions of pH < 8). Therefore, for efficient substrate depletion in the temperature range between 20 and    1 ðM1 s1 Þ ½S  exp 10:3  1100 þ 80 °C, the condition: kðSþSO T 4 Þ h i   4200  should apply. In the presence of chloS2 O2  exp 27:6  8 T ride anions, the term [Cl]  exp(19.3) should be added to the previous equation. TCA was not degraded by light-activated S2 O2 8 at room temperature (David Gara et al., 2007), as expected from k2 < 2  106 M1 s1 (Ross et al., 1998).

V.C. Mora et al. / Chemosphere 75 (2009) 1405–1409

Supported on a detailed kinetic analysis of the experimental results, it was corroborated that TCA showed two distinctive reaction paths when attacked by oxidants and reductive species. Because of these properties of TCA, we have shown that the reaction medium generated by thermally activated peroxydisulfate in the presence of formate, turns highly reductive. Formate is an ideal additive, since after reaction it leaves no organic residue in the reaction media. These properties may be used for the mineralization of refractory to oxidation materials such as perchlorinated hydrocarbons (Gonzalez et al., 2007). However, for the S2 O2 8 =KHCO2 system to be efficient, the dissolved oxygen concentration, [O2], should be 1 1 s Þ > 2 adjusted such that the condition kðS þ CO 2 ÞðM 1010  ½O2 =½S applies for any temperature in the range from 20 to 80 °C. The rate constant kðS þ CO 2 Þ stands for the reaction rate constant between CO 2 and the substrate S. Low oxygen concentrations are self achieved by the reaction system when working with reactors closed to the atmosphere and previous degassing of the reaction system is not necessary. Acknowledgements This work has been supported by the grant PICT 2003 14508 from ANPCyT, Argentina. M.C.G. and J.A.R. are research members of CONICET, Argentina. D.O.M. is a research member of CICPBA, Argentina. V.C.M. thanks CICPBA and CTA YPF, for a graduate studentship. The authors thank Prof. Dr. André M. Braun for the donation of part of the analytical equipment used in this work. Appendix A. Supplementary material Supplementary data associated with this article can be found, in the online version, at doi:10.1016/j.chemosphere.2009.02.038. References Alegre, M.L., Geronés, M., Rosso, J.A., Bertolotti, S.G., Braun, A.M., Mártire, D.O.,  Gonzalez, M.C., 2000. Kinetic study of the reaction of chlorine atoms and Cl2 radicals anions in aqueous solutions. I. Reaction with benzene. J. Phys. Chem. 104, 3117–3125. Anipsitakis, G.P., Dionysiou, D.D., 2004. Radical generation by the interaction of transition metals with common oxidants. Environ. Sci. Technol. 38, 3705–3712. Brusa, M.A., Churio, M.S., Grela, M.A., Bertolotti, S.G., Previtali, C.M., 2000. Reaction volume and reaction enthalpy upon aqueous peroxodisulfate dissociation:  S2 O2 8 to 2SO4 . Phys. Chem. Chem. Phys. 2, 2383–2387. David Gara, P.M., Bosio, G.N., Gonzalez, M.C., Mártire, D.O., 2008. Kinetics of the sulfate radical-mediated photo-oxidation of humic substances. Int. J. Chem. Kinet. 40, 19–24.

1409

David Gara, P.M., Bucharsky, E., Wörner, M., Braun, A.M., Mártire, D.O., Gonzalez, M.C., 2007. Trichloroacetic acid dehalogenation by reductive radicals. Inorg. Chim. Acta 360, 1209–1216. Gonzalez, M.C., Carrillo Le Roux, G., Rosso, J.A., Braun, A.M., 2007. Mineralization of CCl4 by the UVC-photolysis of hydrogen peroxide in the presence of methanol. Chemosphere 69, 1238–1244. Hermann, H., Reese, A., Zellner, R., 1995. Time-resolved UV/VIS diode array absorption spectroscopy of SO x (x = 3, 4, 5) radical anions in aqueous solution. J. Mol. Struct. 348, 183–186. House, D.A., 1962. Kinetics and mechanism of oxidations by peroxydisulfate. Chem. Rev. 62, 185–203. Huang, K.C., Zhao, Z., Hoag, G.E., Dahmani, A., Block, P.A., 2005. Degradation of volatile organic compounds with thermally activated peroxydisulfate oxidation. Chemosphere 61, 551–560. Huie, R.E., Clifton, C.L., 1990. Temperature dependence of the rate constants for reactions of the sulfate radical, SO 4 , with anions. J. Phys. Chem. 94, 8561–8567. Johnson, R.L., Tratnyek, P., Johnson, R.B., 2008. Persulfate persistence under thermal activation conditions. Environ. Sci. Technol. 42, 9350–9356. Killian, P.F., Bruell, C.J., 2003. Thermally activated peroxydisulfate oxidation of polychlorinated biphenyls (PCBs). In: Calabrese, E.J., Kostecki, P.T., Dragun, J. (Eds.), Contaminated Soils, vol. 8. Amherst Scientific Publishing, Amherst, USA, pp. 27–38. Kin, S., Choi, W., 2002. Dual photocatalytic pathways of trichloroacetate degradation on TiO2: effects of nanosized platinum deposits on kinetics and mechanism. Phys. Chem. B 106, 13311–13317. Liang, C., Bruell, C.J., 2008. Thermally activated persulfate oxidation of trichloroethylene: experimental investigation of reaction orders. Ind. Eng. Chem. Res. 47, 2912–2918. Liang, C.J., Bruell, C.J., Marley, M.C., Sperry, K.L., 2003. Thermally activated persulfate oxidation of trichloroethylene (TCE) and 1,1,1-trichloroethane (TCA) in aqueous systems and soil slurries. Soil Sediment Contam. 12, 207–228. Liang, C., Bruell, C.J., Marley, C.M., Sperry, K.L., 2004. Peroxydisulfate oxidation for in situ remediation of TCE. II. Activated by chelated ferrous ion. Chemosphere 55, 1225–1233. Liang, C., Wang, Z.S., Mohanty, N., 2006. Influences of carbonate and chloride ions on peroxydisulfate oxidation of trichloroethylene at 20 °C. Sci. Total Environ. 370, 271–277. Liang, C., Wang, Z.S., Bruell, C.J., 2007. Influence of pH on peroxydisulfate oxidation of TCE at ambient temperatures. Chemosphere 66, 106–113. Madhavan, V., Levanon, H., Neta, P., 1978. Decarboxylation by SO 4 radicals. Radiat. Res. 76, 15–22. Mártire, D.O., Gonzalez, M.C., 2001. Aqueous phase kinetic studies involving intermediates of environmental interest: phosphate radicals and their reactions with substituted benzenes. Prog. React. Kinet. Mec. 26, 201–218. Osgerby, I.T., 2006. ISCO technology overview: do you really understand the chemistry? In: Calabrese, E.J., Kostecki, P.T., Dragun, J. (Eds.), Contaminated Soils, Sediments and Water. Springer, NY, USA, pp. 287–308. Ross, A.B., Mallard, W.G., Helman, W.P., Buxton, G.V., Huie, R.E., Neta, P., 1998. NDRL-NIST Solution Kinetics Database. Ver. 3.0, Gaithersburg, MD, Available on the web at: (March 2008). Waldemer, R., Tratnyek, P.G., Johnson, R., Nurmi, M.T., 2007. Oxidation of chlorinated ethenes by heat-activated persulfate: kinetics and products. Environ. Sci. Technol. 41, 1010–1015. Wardman, P., 1989. Reduction potentials of one-electron couples involving free radicals in aqueous solution. J. Phys. Chem. Ref. Data 18, 1637–1755.  Yu, X.-Y., Bao, Z.-C., Barke, J.R., 2004. Free radical reactions involving Cl, Cl2 , and 2 in the 248 nm photolysis of aqueous solutions containing S O and Cl. J. SO 2 8 4 Phys. Chem. A 108, 295–308.