Environ. Sci. Technol. 1992, 26, 944-951
Dark and Photoassisted Fe3+-Catalyzed Degradation of Chlorophenoxy Herbicides by Hydrogen Peroxide Joseph J. Plgnatello The Connecticut Agricultural Experiment Station, P.O. Box 1106, New Haven, Connecticut 06504
The herbicides 2,4-dichlorophenoxyaceticacid (2,4-D) and 2,4,5-trichlorophenoxyaceticacid (2,4,5-T) were degraded in acidic aerated solutions of H202and Fe2+or Fe3+. Conditions leading to complete mineralization could be achieved using Fe3+/H2O2, which thereby became the focus of the study. Herbicide transformation by Fe3+/H,02 was sensitive to pH (optimum, 2.7-2.8) and was inhibited by methanol or chloride due to scavenging of the active oxidant, and by sulfate due to complexation of Fe3+. The corresponding polychlorophenol was a transient, low-yield intermediate. Dechlorination of 0.1 mM herbicide was rapid and quantitative. Conversion to C02 ranged from about 40 to 70%, depending on [H202](10-500 mM) and was independent of Fe oxidation state. Degradation was markedly accelerated by irradiation with visible light containing a small UV component. Photoassisted conversion of herbicides to COz was quantitative in less than 2 h using H202to herbicide molar ratios as low as 5. The stoichiometry indicated dioxygen consumption.
Introduction Technologies for the destruction or detoxification of hazardous organic wastes are in urgent demand because of growing restrictions placed on land disposal and the need to clean up existing waste sites. Incineration, although generally effective, is not always feasible or economical. Alternatives are needed for small-scale waste, aqueous wastes, and dilute wastes in soil or sludges, where transport off-site is objectionable, or when there are concerns about incinerator emissions ( I , 2). The intent of this study was to explore the use of hydrogen peroxide with iron salts as an alternative oxidant, specifically here for degradation of aqueous pesticide wastes. Hydrogen peroxide in the presence of Fen and Fern produces reactive hydroxyl (HO') radicals and possibly other reactive species (see below). The reagent components are easy to handle and environmentally benign, making this system attractive for treating aqueous or soil-bound contaminants. Pesticide wastes-as on-site container/machinery rinsates, small amounts of unused product, and spills to the soil-have long been a problem for which practical treatment alternatives by both biological and chemical means have been sought (3-7). Historically, Fe/H202oxidations have been of interest mostly from a synthetic/mechanistic perspective and for their relevance to enzyme reactions and free-radical damage to cells. Recent studies have investigated their potential for waste treatment. Usually ferrous iron (Fenton's reagent) has been used (8-12), but some studies have employed ferric iron (13) or ferric oxide (14). The utility of Fe/H202oxidants for organic wastes is still unclear, however, especially because of high peroxide requirements. In addition, little is known about ferric systems. This study concerns the degradation of 2,4-D and 2,4,5-T. The goal was to achieve complete mineralization to eliminate concerns about organic byproducts. 2,4-D and related compounds are important postemergent herbicides. 2,4,5-T and related compounds served a similar purpose until they were discontinued in the United States in the 944
Envlron. Sci. Technol., Vol. 26, No. 5, 1992
mid-19709. However, 2,4,5-T wastes still exist; for example, EPA has large quantities of wastes resulting from cancelled registration that need to be destroyed, but as yet there are no commercial incinerators in the United States licensed to accept them, partly due to public fears of dioxin emissions. Background Chemistry. Both Fen and Fern react with hydrogen peroxide. The "classical" mechanisms for these reactions involve oxyl radical intermediates that can attack organic compounds. The classical reaction of Fe" with H202,known as the Fenton reaction (15, 16), generates OH' in the rate-limiting step (eq 1); O H may be scavenged by reaction with another Fez+(eq 2) or react with an organic compound.
-
Fe2++ H202
Fe3+ + OH'
-
(1)
HO- + Fe3+ (2) Fenl catalytically decomposes H202to O2and H20. The classical "radical chainn mechanism proposed for simple Fe3+(aq)systems (i.e., no complexing ligands other than water) involves OH' and the hydroperoxyl radical (HO,') by the following steps, inclusive of eqs 1 and 2 (15,17,18): Fe3+ + HZO2* Fe-OOH2+ + H+ (3) OH'
+ Fez+
+ HO-
-
Fe-OOH2+
Fe2++ H02' Fe3+ + H02'
OH'
+ Fez+ Fe3+ + H02Fez+ + O2 + H+ HzO + HOz' HOz'
+ HzO2
-+
(4)
(5) (6)
(7)
In the presence of excess peroxide, [Fe2+]is small relative to [Fe3+],since reaction 4 is generally much slower than reaction 1 (17). Reaction 7 is an additional mechanism for OH' scavenging. Hydroxyl reacts rapidly and nonselectively with most organic compounds by H-abstraction and addition to C-C unsaturated bonds (15, 19). Compared to OH', H02' is much less reactive (20), and its conjugate base 02'-(pK,, 4.8) is practically unreactive as a free radical (21). Carbon-centered radicals generated by oxyl radical attack may react with 02,if present, to give organoperoxy radicals (ROO'), which can decompose to form H02' or, ultimately, nonradical oxygenated products (22,23). Organoradicals may in some cases be reduced by Fe" or oxidized by Fe"' (15). Additional pathways for Fe/H202reactions have been introduced that involve other potential oxidants as intermediates. Fenton's reaction was suggested to give a "caged" OH' loosely bound to Fe"' that is capable of intramolecular oxidation of an organoligand (24). Numerous studies have proposed high-valent iron-oxo intermediates, such as Fe03+ (formally FeV) (25) and ferry1 complexes (L)FeIV=O and (L*+)FeIV=O(26-31), where L is an organoligand and L*+is a one-electron oxidized organoligand. A t this point, confirmation of these species in simple noncomplexing systems has been elusive (15,18, %), except possibly in concentrated alkali (32). Some target organic substrates or their degradation intermediates, however,
0013-936X/92/0926-0944$03.00/0
0 1992 American Chemical Society
may form Fe"' complexes that can produce such species. (L)Fe'"=O appears to be a weaker oxidant than OH' (26,28,33). However, note that (L'+)FeTV=Owhere L = porphyin is implicated as the active complex in cytochrome P-450 that carries out epoxidation and aliphatic and aromatic hydroxylations (29). Experimental Section Analyses. 2,4-D, 2,4,5-T, 2,4-dichlorophenol (DCP), and 2,4,5-trichlorophenol (TCP) were analyzed by HPLC on a 25-cm, 5-pm Spherisorb ODs-2 C-18 column (Alltech) using UV detection at 230 nm. The mobile phase (1.5 mL/min) was methanol/water/trifluoroacetic acid (TFA) in the ratio 60:40:0.08 for 2,4-D and DCP and 70:30:0.064 for 2,4,5-T and TCP. Standards (Aldrich, 198% pure) were prepared in 60:40 methanol/water. Samples (2 mL) of the reaction mixture were added to 0.13% TFA in methanol (3 mL), which quenched degradation (see Results). DCP and TCP in CHzClzextracts were confirmed by GC/MS (Hewlett-Packard quadrupole 5988A, full-scan E1 mode) by comparison with authentic samples. Ferrous iron was determined spectrophotometrically as the 2,2'-dipyridyl complex at 520 nm (16) in 10 mM dipyridyl, 1M NH4F, and 1 M sodium acetate. Hydrogen peroxide was determined iodometrically in the presence of fluoride (1OO:lmolar ratio) to prevent interference from Fe3+ (34). Chloride was measured potentiometrically (Orion 96-17B electrode) using NaCl standards. Prior to measurement, 5-mL samples were treated with 0.5 mL of 1M NaHC03 (as buffer), 0.1 mL of Orion ISA reagent, and 0.05 mL of methanol. Carbon-14 was measured by liquid scintillation in 15 mL of Opti-Fluor (Packard Instrument Co.) using the external standard method. Methods. Stock solutions of ferrous and ferric salts (0.1-1 M) were made fresh in 0.1 or 0.01 M of an appropriate mineral acid. Most experiments used Fe(C104)2. 6Hz0 (Aldrich, 98% pure) or Fe(C104)3(GFS Chemicals, Columbus, OH) in HC104. Herbicide stock solutions (0.03-0.14 M) were prepared in water by adding NaOH to a final pH of 6-7. [ring-UJ4C]-2,4-D (10 mCi/mmol), [carbo~y-'~C]-2,4-D (9.0 mCi/mmol), and [ring-U-14C]2,4,5-T (7.1 mCi/mmol) were purchased from Sigma (>98% pure). Experiments were conducted in a climate-controlled room at 21 f 1 "C or in a water bath at 21 f 0.2 "C. A typical experiment with Fe3+ involved preparation of a solution (usually 100 mL in a 250-mL Erlenmeyer flask) containing, as appropriate, herbicide, background electrolytes, and the iron stock solution, which lowered the pH to -3. The pH was then adjusted with acid (usually HC104),or a predetermined amount of NaOH was added and the solution allowed to stand for 15 min. The reaction was initiated by adding HzOz. For the Fez+reactions, the iron and peroxide stock solutions were added in quick succession. The solutions were churned vigorously by a magnetic stirrer to maintain oxygen saturation. Flasks were insulated from heat given off by the stir plate. Experiments in which 14C02was monitored were carried out on a 10-mL scale in 50-mL Erlenmeyer flasks. The flasks were fitted with a stopper holding a polypropylene center well (Kontes) containing a plug of glass wool impregnated with 0.3 mL of 1 M aqueous ethanolamine as COz trap. The flasks were shaken at 100 revolutions per minute. Replicate flasks were removed and sampled for solution-phase I4C, and the center well and its contents (including the glass wool) were added to Opti-Fluor after first separating the glass wool from the well with forceps, Control experiments showed that mass transfer of 14C02 from solution to the well occurred within minutes. Sam-
Table I. Reaction of 2,4-D with Fenton's Reagent (Fe2+/HzO2)"
initial pH
final pH
% loss 2,4-D
1.0 2.0 3.0 4.5 6.0
1.0 2.0 3.0
79 88 76 86 85
3.2 3.4
% yield DCP
34 16 12
7.5 7.5
"Conditions: [Fez+] = [H202] = 0.25 mM; [2,4-D] = 0.1 mM; ionic strength 0.2 M (NaC10.J.
ples for solution I4C (1mL) were mixed with water (1mL) and methanol (0.5 mL) and purged with a stream of argon for 2 min to drive off 14C02before adding scintillation cocktail. Photolysis. Reactions were carried out using 100 mL of solution in identical 250-mL Pyrex Erlenmeyer flasks (Corning 7740 glass, cutoff, -300 nm) covered with a 38mm-diameter watchglass. The solution path length was estimated to be 2.5 cm. Samples under "bright" light were irradiated -30 cm from a rack of four 200-W "cool white" fluorescent lights. Samples under "dim" light were placed further from the source and received both direct and reflected light. Light intensity in the region 400-700 nm was measured with a Li-Cor quantum radiometer (Model LI185B) and in the region 290-385 nm with an Eppley UV radiometer (No. 23520) covered with a Pyrex beaker. Photon intensity was determined by ferrioxalate actinometry at 0.15 M potassium ferrioxalate (35). Results Dark Reaction. (A) Herbicide Transformation by Fe2+/HzOz.The transformat.ion of 2,4-D (0.1 mM) with Fenton's reagent in air-saturated acidic solution was quite fast. When [Fez+]and [H,02] were both greater than 1 mM, 2,4-D was transformed completely in > [Fe3+](17, 25), peroxide decomposition was first order, and the pH dependence of the pseudo-first-order rate constant kObsis given in Figure 4. By comparison of Figure 4 with the 946
3.0 -
Environ. Sci. Technol., Vol. 26, No. 5, 1992
Figure 3. Fe3+/H,O2 mineralization of [rir1g-'~C]-2,4,5-T: [2,4,5-T] = 0.175 mM (1 X lo4 dpm/mL), [Fe3+] = 1.0 mM, [H202] = 10.2 mM, pH 2.75, = 0.2 M (NaCIO,). TCP, 2,4,5-trichlorophenoI.
2,4-D transformation curves in Figure 1, it is seen that these reactions have practically the same pH profile. Reactions were sensitive to the anion in solution originating from background electrolytes and Fe3+ or H+ counterions. The loss of 2,4-D may be approximated as zero order in [2,4-D] (i.e., linear with time) near optimum pH, although more complex kinetics exist at other pH. The pseudo-zero-order rate constant at constant ionic strength (Na salts) followed the order (Table 11) Clod- NO3- > C1- = S042-
Fe3+-catalyzeddecomposition of hydrogen peroxide was also sensitive to the anion; the pseudo-first-order rate
1.o I
Table 111. Extent of Herbicide Mineralization by Dark Fe3+/Hz02at pH 2.7-2.8, g = 0.2 M (NaClO,)
s
labeled compd, mM
0
ring 2,4-D
0.1
Is,
0.1
0
0.1 0.1 0.1 0.1 0.1 0.1 0.1
-1
pH 3.3'
0.1 -..
0
2
1
4
3
5
Hours 1 .o
L
[Fe3+], mM
0 10 100
1
100
0
1 1 1
0.2
53 39 27 56
1
58
5 20
52 57 40 40 66 59
45 56 69 na na na na na na 37 41
500 10 10 10
10
0.5
100 100
carboxy 2,4-D 0.1 ring 2,4,5-T 0.1
10 10
(I
% initial 14Cn remaining in soln as 14C02
[HZOz], mM
1 1 1 1
na, not analyzed.
n
-
c
/
I
E
\
a
0.1
8 e
v)
fi 0
-
Y
0.6
-
0.4
-
0.2
-
P)
3 0.01
I3
1
2
3
4
5
6
7
PH Flgure 4. Fe3+-catalyzeddecomposltlon of hydrogen peroxide: [Fe3+] = 0.99 mM, [H202],= 100 mM, pH 2.8, p = 0.2 M (NaCIO,). Error bar is f 2SD. Note: /cobs incorporates terms which Include [Fe3+] and [H+] .
constant in the absence of 2,4-D followed the order (Table 11) Clod- = NO3- > C1- >> S042A further distinguishing feature of reactions in chloride is the much higher maximum yield of DCP (Table 11). Regardless of the anion, however, DCP was eventually degraded. (C) Herbicide Mineralization by Fe3+/H202. At optimal pH 2.8, Fe3+/H202released chloride rapidly and quantitatively from 2,4-D (Figure 2) and 2,4,5-T (Figure 3). Throughout the course of the reaction, the sum of C1 in herbicide, polychlorophenol, and C1- accounted for between 90 and 100% of total C1 for 2,4-D and between 93 and 115% of total C1 for 2,4,5-T. Therefore, dechlorination was practically concomitant with loss of parent material, and no major amounts of other chlorinated organics can be expected. In reactions of ring-labeled2,4-D, the percent of solution 14C extracted by an equal volume of ethyl acetate declined from 100% at time zero to -5 M. Chloride, on the other hand, inhibits hydrogen peroxide decomposition only slightly (
