Environ. Sci. Technol. 1994, 28,394-400
Chemical Oxidation of Chlorinated Organics by Hydrogen Peroxide in the Presence of Sand Joseph X. Ravikumar and Mirat D. Gurol'
Department of Chemical Engineering and Environmental Studies Institute, Drexel University, Philadelphia, Pennsylvania 19 104 The objective of this study was to investigate the feasibility of using hydrogen peroxide (HzOz) as a chemical oxidant for in-situ treatment of contaminated surface soils. The process has been tested in the presence and absence of ferrous sulfate on sand-packed columns, which contained pentachlorophenol (PCP) and trichloroethylene (TCE) as model compounds. Both column and batch studies have demonstrated that HzOz decomposed readily by interacting with the natural iron content of sand, and additional ferrous salts further enhanced the extent of HzOz decomposition. As a result, PCP and TCE adsorbed on the sand surface were oxidized effectively and by stoichiometric release of organic bound chlorine as chloride ion.
Recently, hydrogen peroxide has been used to facilitate biotreatment of subsurface contamination ( 1 0 , l l ) . These studies have emphasized the fact that hydrogen peroxide might serve as an oxygen source for microorganisms, but have failed to address its capability as a chemical oxidant, which is the focus of our study. Pentachlorophenol (PCP) and trichloroethylene (TCE), which are frequently encountered at many hazardous waste sites as soil and groundwater contaminants, were used in this study as model compounds. Approach
Among the limited number of alternatives available for remediating contaminated surface soils, in-situ remediation techniques are generally more desirable for creating the least exposure to hazardous substances of workers and the general public. However, in-situ techniques such as soil washing, vapor extraction, and steam stripping are used merely for the separation of certain groups of organic contaminants from the soil. As a result, it is often necessary to provide additional treatment for elutriate from soil washing, air emission from vapor extraction, and condensate from steam stripping operations. In-situ bioremediation of unsaturated soils might result in the destruction of certain hazardous organics; however, many compounds, especially those with chlorine substitution, persist in the environment due to their low rates of biodegradation and/or toxicity to microorganisms, among other factors. Therefore, this study was conducted to investigate the feasibility of chemical oxidation as a potential in-situ treatment technology for soils contaminated with chlorinated hydrocarbons. In this study, hydrogen peroxide was selected as the chemical oxidant due to its following attributes: Even though hydrogen peroxide itself is not a good oxidant for most organic substances of interest, it is known to decompose into a highly reactive hydroxyl radical (OH') in the presence of iron salts (1)and supported iron oxides (2). In fact, the mixture of hydrogen peroxide with ferrous salts is known in the literature as Fenton's reagent. The hydroxyl radical, which is believed to be an intermediate product of the interaction of hydrogen peroxide with ferrous and ferric ions and iron oxides, reacts with organic substances rapidly with second-order rate constants in the range of lO7-lOl0M-'s-1 (3). Numerous investigators have found mixtures of hydrogen peroxide and iron salts quite effective in the oxidation of alcohols, ethers, dyes, chlorinated phenols, chlorobenzene, and other chlorinated organics in aqueous solutions and in wastewaters (4-9). As opposed to other oxidants such as ozone, solutions of hydrogen peroxide and iron salts can be readily applied to surface soils for the purpose of in-situ treatment.
In order to simulate in-situ treatment conditions, the main experiments were conducted by using sand columns, which were intermittently applied with solutions of hydrogen peroxide, ferrous sulfate, and distilled/deionized (pure) water. Pure water was used for control experiments in order to observe the desorption of the model compounds from sand. The segmented Teflon columns that were 91 cm in length and 4.0 em in diameter had four sampling ports located 20 cm apart. The columns were acid-washed and sterilized prior to use. Additional experiments were carried out in batch systems by using glass flasks to investigate the interaction of sand with hydrogen peroxide. The properties of the sand (primarily quartz) are presented in Table 1. This coarse-grained sand was selected because the pollution problems are generally worse with free draining granular soil, and the surface soils-the region most affected by an accidental spill-is often predominantly made up of sand. Furthermore, due to its low natural organic composition, this sand was expected to have low interaction with PCP and TCE and thus results in better interpretation of the experimental data. However, it is recognized that the high natural organic content of the soils might affect the applicability of the proposed oxidation technique by (i) binding with the organic compounds, (ii) competing for the chemical oxidant, and (iii) interacting with inorganic soil constituents. In fact, an earlier study (13), in which the oxidation of PCP by Fenton's reagent was investigated in a batch mixed reactor for different types of soil, has reported significantly different PCP oxidation rates for soils with high (0.6%) and low (0.05%) organic carbon content. Prior to each column experiment, a predetermined volume of the solutions of PCP, TCE, or distilled/deionized (pure) water was added to the empty columns. A known amount of the sterilized sand was then aseptically transferred at a uniform rate to the columns. The columns were externally tapped to avoid any dead pockets and to obtain a tight packing. The sand liquid mixture was allowed to stand for a while, after which time the free liquid on top of the added sand was drained. In order to obtain a uniform distribution of sand and the compounds throughout the depth of the columns, this procedure was repeated several times until the columns were completely packed. Peristaltic pumps were used to deliver pure water
384
0013-936X/94/0928-0394$04.50/0
Introduction
Environ. Sci. Technol., Vol. 28, No. 3, 1994
0 1994 American Chemical Society
Table 1. Characterization of Sand8 Used in This Study parameter
value
organic content (wt %) iron content (mg/kg) uniformity coefficient (Cu) coefficient of gradation (Cg) effective particle size ( D d(mm) av particle size diameter (D,,Jb(mm) specific surface areaC(cm2/g) av pore sized (mm) porosity
0.04 800 2.83 1.79 0.24 0.48 30.2 0.048 0.185
"The fraction of the and that was retained on sieve no. 18 (corresponds to 1mm size). * Estimated from the expression ( E M ) / (ZMiDi),where Mi is the mass of sand retained between two adjacent sieves and Di is the mean diameter of the adjacent sieves (12). c Calculated assuming spherical-shaped particles and a soil particle density of 2650 kg/m3. For column application, calculated from the expression Dl0/5. e For column application, determined from the expression [l - (sand bulk densityhand particle density)].
or solutions of hydrogen peroxide and iron salt intermittently to the columns. Silicon tubing was used to connect the pumps to the columns. All other tubings and fittings were of Teflon or glass. The experimental setup is illustrated in Figure 1. The study was conducted in three phases. The first phase addressed the decomposition of hydrogen peroxide over the sand columns as well as in batch systems in the absence of PCP and TCE. For column experiments, a solution of hydrogen peroxide (170 mg/L) was applied to one of the sterile sand columns while hydrogen peroxide and a ferrous sulfate solution were applied simultaneously to the second one by using two separate pumps. A total of 100 mL (by volume) of ferrous sulfate a t 50 mg/L concentration was applied to the column. Samples were collected from the sampling ports of the columns a t various time intervals and were analyzed for hydrogen peroxide concentration. The batch experiments were conducted in order to separate and identify the variables responsible for the decomposition of hydrogen peroxide, e.g., iron content, surface area, and organic content of the sand. The effect of surface area on the decomposition of hydrogen peroxide was studied by using glass beads. Glass has chemical characteristics similar to that of sand, except that it does not contain any measurable iron. The glass beads used in the experiments were 1 mm in diameter, which is identical to the size of the sand grains used in the study. Additional experiments were conducted with sand of different sizes (0.5 and 1.4 mm) but with the same iron and natural organic content. Furthermore, in order to ascertain the effect of the natural organic content of sand on the decomposition of hydrogen peroxide, additional experiments were conducted with sand that was heated in a muffle furnace a t 550 "C for about 1 h to eliminate its organic matter. The batch experiments were conducted in wide-bottomed Erlenmeyer flasks of 1000-mL capacity under sterile conditions. Glass or sand was contacted with 100 mL of hydrogen peroxide solution at 175 mg/L concentration. A known volume of ferrous sulfate solution was added to three of the four flasks which contained glass beads. No ferrous sulfate was added to those flasks which contained sand. The flasks were covered with parafilm and were shaken for about 1 h on a shaker in order to improve the contact between the solution and the glass or sand. Solutions were then analyzed for hydrogen peroxide
concentration. A total of 27 experiments were conducted, including the control experiments. The mass of glass or sand used in the experiments and the corresponding surface area are presented in Table 2. The second phase of the study focused on the removal of PCP and TCE from the sand columns by hydrogen peroxide. Pure water was applied to the control columns at the same rate as hydrogen peroxide solutions to observe the desorption of PCP and TCE from the columns. The operating characteristics of the columns are presented in Table 3. Aqueous samples (leachates) were collected from the sampling ports of the columns and were analyzed for PCP or TCE. The initial and the residual concentrations of the model compounds in sand were determined following extraction of PCP by a solution of sodium hydroxide (14) and extraction of TCE by hexane. Mass balances for the compounds were performed based on the measurements of the initial and the residual concentrations in sand and the concentrations measured in the leachates. The final phase of the study addressed the additional effect of ferrous sulfate on the removal of PCP and TCE from the sand columns. For comparison purposes, one column was treated with a mixture of hydrogen peroxide (100 mg/L, 40 mL) and ferrous sulfate (50 mg/L, 20 mL), a second one was treated with a solution of hydrogen peroxide (100 mg/L, 60 mL), and the control column was treated with pure water (60 mL). The operational characteristics of the columns were the same as thosegiven in Table 3. Aqueous samples were collected from the sampling ports of the columns and were analyzed for PCP or TCE, the UV absorption of the samples a t 254 nm, and the concentration of chloride ion. Measurement of chloride ion provided information on the extent of chloride release due to the cleavage of the C-Cl bonds during oxidation, while the UV absorption served as a measure of the aromatic stability of PCP against chemical oxidation. In homogeneous solutions of hydrogen peroxide and iron salts, the pH was observed to have a significant effect on the decomposition of hydrogen peroxide and the oxidation efficiency of organic compounds. This pH dependence was attributed to the speciation of iron (1)as well as the acidity of the peroxy radical (6),the free radical which is believed to propagate the chain reaction. The pH was not studied as a variable in this feasibility study because of the assumed impracticality of using acidic/basic solutions in the field for in-situ applications. During our experiments, the pH of the solutions remained relatively constant within the pH range of 6-7.
Analytical Methods Hydrogen peroxide concentration was measured by the iodometric method (15) in the absence of any organic matter. A Milton Roy Spectronic 1001 split-beam spectrophotometer was used at a wavelength of 254 nm to measure the UV absorption of aqueous solutions. Chloride ion concentration was measured with an Orion specificion electrode after sample dilution. The concentration of PCP was measured by a Varian high-pressure liquid chromatograph (HPLC) equipped with a UV detector (254 nm). A mixture of 55% acetonitrile/45% water was used as the mobile phase on a Varian MicroPak MCH-10 reverse-phase column. The concentration of TCE was measured by a Varian gas chromatograph (GC) equipped with a Tekmar purge-and-trap concentrator and flame Environ. Sci. Technol., Vol. 28, No. 3, 1994
395
Soil- induced decomposition of Hydrogen Peroxide
DesorpQn expeliment
Chemicaloxidation experiment
soil
soil
and
'and PCP or TCE
PCP or TCE
h
,::$ Poll c
Port D
Figure 1. Experimental setup for column studies. ~
Table 2. Mass of Glass o r Sand a n d t h e Corresponding Surface Area type of material
size of material (mm)
specific surface area (cm2/gP
mass of material (g)
total surface area (cm9
glass glass glass sand sand sand sand sand sand sand sand sand
1.0 1.0 1.0 0.5 0.5 0.5 1.0 1.0 1.0 1.4 1.4 1.4
22.6 22.6 22.6 45.3 45.3 45.3 22.6 22.6 22.6 16.2 16.2 16.2
5.0 7.5 10.0 5.0 7.5 10.0 5.0 7.5
113 170 226 226 340 453 113 170 226 81 121 162
10.0 5.0
7.5 10.0
Specific surface area calculated assuming spherical size particles and a density of 2650 kgim3. Table 3. Operational Characteristics of S a n d Columns descriptor
values
vol of solution delivered per application (mL) rate of solution application (mL/min) superficial velocity (cm3 cm-2 min-l) av linear velocity (cm3 cm-2 min-l) pore vol (mL) empty bed contact time (rnin) seepage time (min) total no. of applications no. of pore volume for total application
60 or 180
0
18
1.4 7.7 200 61 11 4 O or 15b
3.6O or 4.5O
Corresponds to 180 mL. Corresponds to 60 mL.
ionization detector. The column was 8 ft X 1/8 in. 0.d. stainless steel column packed with 1% SP-1000 on 60/80 Carbopack B. The carrier gas was helium a t a flow rate of 40 mL/min. The column temperature was initially set at 45 "C for 3 min, was increased to 220 "C at 8 "Cimin, and then held at 220 "C for 15 min. The injector and the detector temperature were 200 and 250 "C, respectively. For measurement of the PCP and TCE'concentrations in the sand, the sand columns were dismantled and the contents were transferred to and homogenized in a large beaker. For PCP measurements, 20 g of the homogenized sand was mixed in another beaker with 20 mL of 0.1 N 396
Environ. Sci. Technol., Vol. 28, No. 3, 1994
sodium hydroxide. The mixture was allowed to stand for 15 min, after which time the liquid layer was decanted to a tube and centrifuged a t 2000 revolution/min for 5 min. The supernatant was decanted and filtered with the aid of a Swinny syringe unit. The filtrate was acidified to pH 3 with glacial acetic acid, and the sample was injected into HPLC (16). For measurement of TCE, a mixture of hexane and sand prepared at a 1:3 ratio (w/w) was vortexed and placed in a sonicator bath at room temperature for 1 h. This mixture was centrifuged for 30 min at 5000 rpm, and the TCE in the resultant supernatant was measured on a GC. The iron content of the soil was determined in accordance with the following procedure. A 250-mL sample of a 0.25 M solution of hydroxylamine hydrochloride in concentrated hydrochloric acid was heated in a 500-mL glass bottle to 50 "C in a hot water bath. A total of 400 mg of the sand was added to the warm solution, and the sample was returned to the hot water bath (50 "C) and gently shaken for 60 min. The sample was removed from the water bath after 60 min and centrifuged at 10000 revolution/min for 15 min (17). The concentrate was then analyzed for total iron on a Perkin-Elmer atomic absorption spectrophotometer with a flame furnace at a wavelength of 248.3 nm and a slit width of 0.2 nm. Finally, the soil organic carbon was determined by the wet combustion method (18).
Results and Discussion Decomposition of Hydrogen Peroxide. Solutions of hydrogen peroxide were applied t o the sand columns which contained no PCP or TCE, and the concentration of hydrogen peroxide was measured in samples collected from the sampling ports along the depth of the columns. According to the results presented in Figure 2 , the concentration of hydrogen peroxide progressively decreased as a function of the depth of the column, and about 60% of the hydrogen peroxide added to the column has decomposed over a 90-cm-deep sand. The low natural organic content of the sand was not expected to be responsible for the consumption of hydrogen peroxide. However, the iron content of the sand, which is 800 mgi kg, was believed to catalyze the decomposition of hydrogen peroxide, as was observed by others for ferrous and ferric ions (1) and iron oxides (2). As shown in Figure 2 , the
.-C
Iron content of sand is 800 m@g.
3
0
6
18 24 Depth (in inches)
12
30
36
Figure 2. Decomposition of H202through the sand columns in the presence and absence of FeS04.
application to the sand columns of external iron as ferrous sulfate along with hydrogen peroxide resulted in far greater decomposition of hydrogen peroxide. This was expected due to a higher amount of iron being in contact with hydrogen peroxide. The results further implied that the soluble form of iron was more effective than the natural iron content of sand per milligram basis in catalyzing the decomposition of hydrogen peroxide, since the addition of 4 mg of soluble iron to 1920 mg of sand-bound iron has increased the decomposition of hydrogen peroxide from 60 to 88%. This could be due to the limited availability of natural iron for the peroxide reaction. No attempt has been made in this study to understand how iron was distributed in the sand particles. Batch experiments were conducted in order to provide conclusive evidence whether the iron content, the surface, or the organic content of the sand was responsible for the decomposition of hydrogen peroxide. The results summarized in Table 4 clearly show that no appreciable decomposition of hydrogen peroxide was observed over the glass surface (experiments 2-4). However, the application of ferrous sulfate along with hydrogen peroxide increased the decomposition of hydrogen peroxide over glass, and the rate of decomposition of hydrogen peroxide was directly proportional to the amount of iron in the solution. In other words, the decomposition of hydrogen peroxide, which was about 4.3% in the presence of 1mg of iron, has increased to about 23% in the presence of 5 mg of iron and to 45% in the presence of 10 mg of iron. The decomposition of hydrogen peroxide computed as percentage per milligram of iron added was found relatively constant (about 4 % per mg iron) over the entire range of iron and the surface area of glass that was examined. In contrast, the decomposition of hydrogen peroxide over sand increased with increasing mass of sand (experiments 13-15; 16-18; 19-21). This is in accordance with expectations, since as indicated earlier, increasing the mass of sand increases the amount of iron as well as the surface area in contact with hydrogen peroxide. The question regarding whether it is the iron content or the surface area that is responsible for the decomposition of hydrogen peroxide was addressed by comparing the decomposition of hydrogen peroxide for different sizes of sand, i.e., different surface area, while keeping the mass of sand identical. Since the amount of hydrogen peroxide decomposition did not show any dependence on the size of the sand grains (experiments 13, 16,19; 14, 17,20; 15, 18, 21), it was concluded that the surface area had no impact on the decomposition of hydrogen peroxide and that it
was the natural iron in sand which was responsible for the decomposition of hydrogen peroxide. Furthermore, the amount of hydrogen peroxide decomposition was measured as about 2.7 % per mg of natural iron for the entire range of particle sizes and masses of sand examined. This confirmed the observation made during column studies that the natural iron of sand was less effective than ferrous sulfate in solution (per mg basis) in decomposing hydrogen peroxide. Since the sand used in this study contained some organic matter (albeit a low 0.04%), additional experiments were conducted with sand containing no organic matter in order to check its possible contribution to the removal of hydrogen peroxide. Table 4 reveals the same results for experiments 22-24 as for experiments 16-18, indicating that the low-level organic matter present in the sand did not have any significant effect on the decomposition of hydrogen peroxide. Treatment with Hydrogen Peroxide. This phase of the column studies has focused on the removal of PCP and TCE from sand by hydrogen peroxide. The concentration of TCE measured in the leachates collected from the columns is presented in Figure 3. It is clear that the application of pure water to the column caused TCE to desorb from the sand rather readily. Similar results were obtained for PCP as well. It is well-known that the binding capacity of an organic compound to the soil is generally a direct function of the soil organic content. Since the sand used in this study had a low organic content, PCP and TCE might have been bound to the sand surface weakly and, hence, leached rather readily from the columns upon application of pure water. On the other hand, the application of hydrogen peroxide to the columns resulted in a substantial decrease in the concentrations of TCE (as well as PCP) in the leachates. We attribute the difference in the concentrations between the control columns (desorption) and the columns treated with hydrogen peroxide to the chemical oxidation of PCP and TCE. It is conceivable that decomposition of hydrogen peroxide over the sand surface might generate hydroxyl radicals through a mechanism similar to those described in the literature for ferrous and ferric iron and that these radicals might be responsible for the oxidation of PCP and TCE in the sand columns. Mass balances for PCP and TCE were performed by measuring the concentrations of these compounds in leachates and in sand after dismantling the columns. In Table 5, the percent removal of the compounds attributable to desorption is compared to percent removal attributable to chemical oxidation. The application of 1.68 mmol of hydrogen peroxide to the sand column resulted in the oxidation of 61% of PCP in the column which initially contained 0.53 mmol of PCP. The same amount of hydrogen peroxide caused 50 % oxidation of TCE in the column, which initially contained 1.10 mmol of TCE. According to these figures, the hydrogen peroxide requirement was calculated as 5.2 mol/mol of PCP oxidized and as 3.1 mol/mol of TCE oxidized. Hence, hydrogen peroxide was found to be more effective on a molar basis in removing TCE than PCP; this observation is in agreement with the oxidation chemistry which involves more intermediate steps and intermediate products in the oxidation of PCP than in the oxidation of TCE. It is expected that the intermediates of PCP oxidation, still having aromatic and olefinic structures, e.g., quinones and Environ. Scl. Technol., Vol. 28, No. 3, 1994
307
Table 4. Decomposition of Hydrogen Peroxide over Glass and Sand Surfaces type of material
mass of material (9)
exp. no.
surface area of material (cm?
iron (mg)
HzOz decompn ( % )
control (no sand or glass)
0.11 0.17 0.17 1.89 1.92 1.89 av = 1.9
1
glass (1.0 mm)
2 3 4
5.0 7.5 10.0
113 170 226
glass (1.0 mm)
5 6 7
5.0 7.5 10.0
113 170 226
1.0" 1.00
glass (1.0 mm)
8 9 10
5.0 7.5 10.0
113 170 226
5.0" 5.0" 5.0"
glass (1.0 mm)
11
12
5.0 7.5 10.0
113 170 226
10.0" 10.0" 10.0"
13 14 15
5.0 7.5 10.0
226 340 453
4.0b 6.0b 8.0b
4.3 4.4 4.2 av = 4.3 21.4 22.6 23.8 av = 22.6 44.8 45.1 45.4 av = 45.1 10.7 16.2 21.6
sand (1.0 mm)
16 17 18
5.0 7.5 10.0
113 170 226
4.0b 6.0b 8.0b
10.8 16.3 21.7
sand (1.4 mm)
19 20 21
5.0 7.5 10.0
81 121 162
4.0b 6.0b 8.0b
10.6 16.4 21.7
22
5.0 7.5
113 170 226
4.0b 6.0b 8.0b
11.0 16.0 22.0
sand (0.5 mm)
sandc (1.0 mm; no organic)
23 24 a
10.0
HzOz (%) decompn per mg of iron
1.0"
4.3 4.4 4.2 av = 4.3 4.3 4.5 4.8 av = 4.5 4.5 4.5 4.5 av = 4.5 2.7 2.7 2.7 av = 2.7 2.7 2.7 2.7 av = 2.7 2.7 2.7 2.7 av = 2.7 2.8 2.7 2.6 av = 2.7
Ferrous sulfate added to provide iron. Iron contributed from the sand. Organics burnt in a muffle furnace.
.
I
Table 5. Mass Balance for PCP and TCE over Sand Columnsa desorption chemical oxidation PCP TCE PCP TCE % of initial quantity of compd in leachate % of initial quantity of compd retained on sand % of compound accountable % of compound unaccountable % of compound oxidized
1
2
3
4
Number of Applications Flgure 3. TCE concentration in leachates from columns treated with water and a solution of H2O2. Column 1, desorption (H20); column 2, chemical oxidation (H202). Volume of pure water or H202 per application, 180 ml; H202per application, 0.42 mmol. TCE in soil, 1.10
mmol.
tetrachloro- and trichlorophenols (19), will continue to exert an oxidant demand and, hence, will compete with PCP for the oxidant. On the other hand, the major stable intermediate product of TCE is formic acid, which upon reaction with the hydroxyl radical is oxidized to the end products, carbon dioxide and water (20). The effect of hydrogen peroxide dosage on the removal of PCP from the sand columns is demonstrated in Table 398
Environ. Sci. Technol., Vol. 28, No. 3, 1994
51.0
56.7
25.5
32.0
44.0
38.3
8.2
13.0
95.0 5.0
95.0 5.0
33.7 66.3 61.3
45.0 55.0 50.0
a Initial PCP on sand = 0.53 mmol. Initial TCE on sand = 1.10 mmol. HzOz applied = 1.68 mmol. Molar ratio for PCP (H202/ initialPCP) = 3.17. Molar ratio for TCE (HzOz/initialTCE) = 1.53.
6. It is evident that the application of higher dosages of hydrogen peroxide causes a lower concentration of PCP in the leachate as well as a lower residual of PCP on the sand. As a result, the amount of PCP removed from the columns increases with increasing dosage of hydrogen peroxide. Treatment with a Mixture of Hydrogen Peroxide and Ferrous Sulfate. Since the introduction of ferrous sulfate to the columns along with hydrogen peroxide caused a significant enhancement in the decomposition of hydrogen peroxide (Figure 2), this mixture was applied to a
Table 6. Effect of HzOz Dosage on PCP Removal applied HzOz dosage (mmol) 0 1.68 2.44
PCP in leachate PCP retained on sand PCP unaccountable
0.270 0.233 0.026
0.135 0.043 0.351
n n,
0.093 0.029 0.407
I
"."J
1
3
5 7 9 11 Number of Applications
13
15
I
0.00 0
1
2
3
Applied H202 (Cumulative) (rnrnoles)
Figure 4. Comparison of PCP concentrationin leachates from columns treated with H202alone and a mixture of H202and FeS04. Initial PCP on sand, 0.53 mmol. Case 1: H202concentration, 100 mg/L; volume per application, 60 mL. Case 2: H202concentration, 67 mglL, and FeS04 Concentration, 16 mg/L; volume per application, 60 mL.
PCP column with the expectation that PCP removal can further be improved. According to Figure 4, in which the PCP concentration in the leachate was plotted as a function of the hydrogen peroxide dosage, the mixture of ferrous sulfate and hydrogen peroxide was indeed more effective than hydrogen peroxide alone in oxidizing PCP. It is also clear that every incremental addition of hydrogen peroxide has removed less PCP on a per mole basis. This is in accordance with the reaction mechanism involving the hydroxyl radical, since the radical, which is generally a nonselective oxidant, would initially oxidize PCP-the predominant organic species in the system. However, as the reaction proceeds, and consequently the PCP concentration is reduced, the reaction intermediates and the soil organic matter would more successfully compete with PCP for the hydroxyl radical, causing a reduction in the oxidation rate of PCP. In Table 6, the mass balances for PCP are compared for treatment with hydrogen peroxide alone and for treatment with the mixture of hydrogen peroxide with ferrous sulfate. The results indicate a hydrogen peroxide requirement of 6.85 mol/mol of PCP oxidized when hydrogen peroxide alone is used as the oxidant. For the mixture, the hydrogen peroxide requirement is reduced to 4.31 mol/mol of PCP oxidized. The lower oxidant requirement for the mixture is an indication of the more efficient and accelerated decomposition of hydrogen peroxide to the hydroxyl radicals in the presence of ferrous sulfate, as elaborated on earlier. The UV absorption of aqueous PCP samples was used for further assessment of the oxidative capability of hydrogen peroxide in the absence and presence of additional iron. The UV absorption of the leachates collected from the columns is plotted in Figure 5. Compared to the leachate from the control column, hydrogen peroxide and the mixture of hydrogen peroxide and ferrous sulfate
Flgure 5. UV absorbance (254 nm) of leachates from PCP columns. Column 1: desorption (60 mL of H20dellveredper application). Column 2: chemical oxidation by H202(60 mL of solution containing 100 mg/L of H202delivered per application). Column 3: chemical Oxidation by H2024- FeS04 (60 mL of solution containing 67 mg/L cf H202and 16 mg/L of FeS04 delivered per application).
exhibited significantly reduced UV absorption, especially for the later stages of the oxidation. A plausible explanation for this phenomenon is that during the early stages of oxidation of PCP most of the initial intermediates, which still possess the aromatic character, will continue to absorb light in the UV range. However, prolonged oxidation would lead to the opening of the aromatic ring and the formation of simpler products, e.g., carboxylic acids, resulting in an overall reduction in UV absorption of the reaction mixture. Furthermore, the observation of even lower UV absorption of the leachate from the column treated with the mixture confirms our contention that hydrogen peroxide in the presence of ferrous sulfate becomes more effective in destroying the aromatic structure and, as a result, forming simpler reaction products. The cumulative amount of chloride liberated as a result of oxidation of PCP is presented in Figure 6. Again the mixture of hydrogen peroxide with ferrous sulfate was found to be more effective than hydrogen peroxide alone in liberating chloride on a per mole hydrogen peroxide basis. Similar results were obtained for TCE as well,
Summary and Conclusions Hydrogen peroxide was investigated as a potential chemical oxidant for the in-situ treatment of soil contaminants. Under the experimental conditions and for the sand used in this study, hydrogen peroxide effectively oxidized PCP and TCE in sand columns both in the presence and in the absence of externally added iron. The results in terms of hydrogen peroxide requirement and the amount of chloride released upon oxidation of PCP and TCE are summarized in Table 7. Hydrogen peroxide in the presence of ferrous sulfate produced consistently better results in terms of lower oxidant requirement and more extensive oxidation of PCP and TCE. The results also demonstrated that TCE exhibited about half as much of the hydrogen peroxide requirement as PCP in the presence and absence of ferrous sulfate. Under the experimental conditions used in this study, chemical oxidation has caused 80-90 76 of organic bound chlorine Environ. Scl. Technol., Vol. 28, No. 3, 1994
388
in the heterogeneous system should first address the reaction mechanisms and kinetics in homogeneous systems, which are currently under investigation in our laboratories. Acknowledgments
The work described here is a part of a study which was financially supported by the U.S.EPA, Office of Research and Development, Superfund Research Program (R815266-0). J.X.R. was a graduate student at Drexel University a t the time of this study. H202alone
Literature Cited
H202cFeS04
0
1 2 Applied HZOZ (Cumulative) (mmoles)
3
Figure 6. Chloride ion measured in leachates from PCP columns.
Table 7. Summary of Oxidant Requirement and Chloride Release treatment oxidant requirement (mol of HzOz consumed/ mol of compd oxidized) chloride release (mol/mol of HzOz consumed) theoretical chloride release (mol/mol of compound oxidized) chloride release (molimol of compd oxidized) chloride release (percent of theoretical)
model compds PCP TCE
HzO2 alone H ~ 0 2+ FeS04
6.9 4.3
3.7 2.1
Hz02 alone Hz02 + FeS04
0.6 1.1
0.7 1.3
HzOz alone Hz02 + FeS04 HzOz alone HzOz + FeS04
5.0 4.1 4.6 02 92
3.0 2.6 2.7 87 90
to be released as chloride ion for both PCP and TCE. This indicates that complete mineralization of such chlorinated compounds might be achievable if a sufficient amount of hydrogen peroxide is used and all the process parameters are identified and controlled properly. Even though the results presented here are quite encouraging for the proposed process, the quantitative resultsare limited to the type of sand and the experimental conditions used in this study. Hence, additional and more extensive testing on different types of soil are needed in order to address the effects of soil characteristics on the proposed process. Furthermore, more fundamental work is required to understand the reaction mechanisms in this heterogeneous system, including the role of iron and iron oxides, surface chemistry, adsorption/desorption characteristics of reactants, complexation of contaminants with soil organic matter, scavenging of hydroxyl radical by soil organic matter, and oxidation reactions at the surface and in solution phase. However, even in homogeneous solutions of hydrogen peroxide and iron salts, the reaction mechanism and the factors affecting the formation and consumption rates of the hydroxyl radical are poorly understood. Hence, any attempt toward solving the puzzle
400
Envlron. Sci. Technol., Vol. 28, No. 3, 1994
(1) Barb, W. G.; Baxendale, J. H.; George, P.; Hargrave, K. R. Trans. Faraday SOC.1959, 47, 591. (2) Kitajima, N.; Matsumura, T.; Fukuzumi, S. I.; Ono, Y. J . Phys. Chem. 1977,81, 1307. (3) Dorfman,L. F.; Adams, G. E. National Bureau of Standards Report No. NSRDS-NBS-46. Government Printing Office: Washington, DC, 1973. (4) Barbeni, M.; Minero, C.; Pelizzetti, E.; Borgarello, E.; Serpone, N. Chemosphere 1987,16,2225. (5) Bishop, D. F.; Stern, G. F.; Fleischman, M.; Marshall, L. S. Znd. Eng. Chem. Process Des. Dev. 1968, 7, 111. (6) Walling, C. Acc. Chem. Res. 1975, 8 , 125. ( 7 ) Bowers, A. R.; Gaddipati, P.; Eckenfelder, W. W., Jr.; Monsen, R. M. Water Sci. Technol. 1989, 21, 477. (8) Carberry, J. B.; Benzing, T. M. Water Sci. Technol. 1991, 23, 367. (9) Kuo, W. G. Water Res. 1992, 26, 881. (10) Spain, J. C.; Milligan, J. D.; Downey, D. C.; Slaughter, J. K. Ground Water 1989,27, 163. (11) Morgan, P.; Watkinson, R. J. Water Res. 1992,26, 73. (12) Das, B. M. Principles of Geotechnical Engineering; PWSKent Publishing Co.: Boston, 1990; p p 63-67. (13) Watts,R. J.;Udell,M.D.;Rauch,P.A.;Leung,S. W. Hazard. Waste Hazard. Mater. 1990, 7, 335. (14) Dao, T. H.; Lavy, T. L.; Dragun, J. Residue Rev. 1983,87, 91. (15) Kolthoff, I. M.; Sandell, E. B. Textbook of Quantitative Inorganic Analysis, 3rd ed.; Macmillan Co.: New York, 1952; p 600. (16) Ravikumar, J. X. Ph.D. Dissertation, Drexel University, Philadelphia, 1992. (17) Hesse, P. R. A Textbook of Soil Chemical Analysis;Chemical Publishing Co., Inc.: New York, 1971; p 332. (18) Nelson, D. W.; Sommers, L. E. In Methods ofsoil Analysis, Part ZI Chemical and MicrobiologicalProperties,2nded.; Page, A. L., Ed.; American Society of Agronomy, Inc., and Soil Science Society of America, Inc.: Madison, WI, 1982; Chapter 29. (19) Wong, A. S.; Crosby, D. G. Water. J. Agric. Food Chem. 1981, 29, 125. (20) Cooper, W. J.; Nickelsen, M. G.; Waite, T. D.; Kurucz, C. N. In Proceedings of A symposium on Advanced Oxidation
Process for the Treatment of Contaminated Water and Air, Toronto, Canada, 1990; Environment Canada, Wastewater Technology Centre: Toronto, Canada, 1990;p p 1-20.
Received for review M a y 3,1993. Revised manuscript received November I , 1993. Accepted November 4, 1993." Abstract published in Advance ACS Abstracts, December 15, 1993.
