Hazardous and Industrial Waste Proceedings, 30th Mid-Atlantic Conference
Hazardous and Industrial Waste Proceedings, 30th Mid-Atlantic Conference
Edition 1st Edition
First Published 1998
eBook Published 17 July 2014
Pub. location Boca Raton
Imprint CRC Press
Pages 939 pages
eBook ISBN 9780429080753
SubjectsEngineering & Technology
Christensen L. (1998). Hazardous and Industrial Waste Proceedings, 30th Mid-Atlantic Conference. Boca Raton: CRC Press, https://doi.org/10.1201/9781498709453
This book is based on the Mid-Atlantic Industrial and Hazardous Waste Conference to bring together professionals interested in the advancement and application of technologies and methods for managing industrial and hazardous wastes.
TABLE OF CONTENTS
MANNERS Manners are the social mores of our society. Within each society we agree on what are good and bad manners. In many societies, when someone holds open a door for us, for example, we respond with a "thank you", acknowledging an unselfish deed. Manners is what the African-
not adhere to a adhere to the admonition will make them better engineers. moral statement is that engineering involves the public and not keeping up must do to continue to be part of the engineering community. The can result in harm to both the engineer and the public.
admonitions statement in a Code of Ethics will not get engineers in trouble, but to For example, in the ASCE Code of Ethics, Guideline 7a reads: 7a. Engineers should keep current in their specialty field by engaging in professional practice, participating in continuing
way. That is, if we find that having bad manners, or acting immorally, or even breaking a law is advantageous to us individually, why we, at any given moment, ought not to want others to emulate. The obvious answer to that question is that we don’t want to get
such actions. But if, by telling lies, I become a scoundrel, why would I believe this to be an undesirable result? If I ran an engineering practice where I made it a habit of telling lies to clients (e.g. "The report is in the mail" when in fact it is still being prepared), why is this detrimental? Why will telling lies
good engineer. Lying: Moral Choice in Public and Private Life On Being Responsible University Press of
Bok, Sissela (1978) Pantheon, New York Pritchard, Michael S. (1991) Kansas, Lawrence KS
P’ A 2 P: P , A 2
DECHLORINATION OF CARBON TETRACHLORIDE BY NANOSCALE IRON PARTICLES IN AQUEOUS SOLUTION HSING-LUNG LIEN WEI-XIAN ZHANG Department of Civil and Environmental Engineering, Lehigh University, Bethlehem, PA 18015 INTRODUCTION Recently, a method for the generation of very small (nanoscale) bimetallic particles has been reported [1,2]. These nanoscale metal particles typically have a diameter on the order of 1-100 nm and feature 0.06% by weight of palladium deposited on the surface of iron. Advantages of the nanoscale bimetallic system for treatment of chlorinated organic pollutants include: (1) High specific surface area. The nanoscale metal particles have a specific surface area around 35 m2/g. Tens to hundreds times higher than those of the commercial grade iron particles (used in conventional iron walls). (2) High surface reactivity. For example, values of surface-area-normalized rate coefficient (KSA) for the transformation of chlorinated ethylenes were about one to two-orders of magnitude higher than those reported in the literature for commercial grade iron particles . Due to their small particle size and high reactivity, the nanoscale metal particles may be useful in a wide array of environmental applications. In the aqueous phase, the nanoscale iron particles remain suspended, almost like a homogenous solution. Theoretical calculations indicate that, for colloidal particles less than about 1 micrometer, gravity of the metal particles is insignificant to influence the particle movement. Brownian motion (thermal movement) tends to dominate the transport process in groundwater. Thus, we believe that the metal particles could be injected directly into contaminated soils, sediments and aquifers for in situ remediation of chlorinated hydrocarbons, offering a cost-effective alternative to such conventional technologies as pump-and-treat, air sparging or even conventional iron reactive walls. Design, construction and operation of such injectable systems should be reasonably straightforward.
We present here applications of the nanoscale metal particles for transformation of carbon tetrachloride (CT). CT is one of the most prevalent contaminants in soils and aquifers. It has been listed as priority pollutants by the U.S. Environmental Protection Agency, and also appeared on the Superfund National Priority List. There is an urgent need to develop effective control and treatment methods. The purpose of this study was aimed to measure the rate and extent of dechlorination, characterize and quantify reaction intermediates and final products. EXPERIMENTAL METHODS Synthesis of nanoscale iron particles. Nanoscale iron particles were synthesized by adding 1:1 volume ratio of FeCI3*6H20 (0.045M) into NaBH4 (0.25M) solution and mixed vigorously under room temperature (22±1 °С) for a few minutes. Ferric iron (Fe3+) was reduced to zero-valent iron (Fe°) by borohydride, a strong reductant. Metal particles from this reaction have sized mostly in the range between 1 to 100 nanometers [1,2]. BET analysis gave a specific surface area of 35 m2/g. Batch experiments. Batch experiments were conducted with 50 mL serum bottles. In each batch bottle, 20 mL deionized water was mixed with 0.25 g of the nanoscale metal particles. Then, 10 pL stock solution of CT dissolved in methanol was spiked into the solution. Initial organic concentration was about 0.1 mM. The serum bottles were capped with Teflon Mininert valves and mixed on a rotary shaker (30 rpm) at room temperature (22±1°C). Parallel experiments were also performed without the metal particles (control) and with a commercial grade iron (Aldrich, 99%, <10 pm, BET surface area 0.9m2/g ). Methods of Analyses. Organic concentrations were measured by the static headspace gas chromatograph (GC) method. At selected time intervals, 20 pL headspace aliquot was withdrawn from the batch bottle for GC analyses. Concentrations of chlorinated methanes were measured using a HP5890 GC equipped with a DB-624 capillary column (30mx0.32mm) and an electron capture detector (ECD). The detection limit of this method was less than 5 pg/L. Hydrocarbon products in the headspace were qualitatively identified with a Shimadzu QP5000 GC-MS and further quantified with GC analysis by comparing retention times and peak areas with standard gas samples (ethane, ethylene, acetylene, methane and carbon dioxide).
RE SULTS AND DISCUSSION Reactions with nanoscale Fe particles. Reactions of CT and CF (chloroform) with nanoscale Fe particles are shown in Figure 1. 0.103 mM (15.86 mg/L) of CT was completely reduced within 20 hours. Several intermediates and final products were observed. CF emerged very quickly and peaked (61%) around 20 hours corresponding to the disappearance of CT. Complete reduction of CF was observed around 100 hours. Both dichloromethane (DCM) (51%) and methane (41%) accumulated steadily and constituted the two major final products. Reactions with commercial grade Fe particles. Reactions of CT with the Aldrich Fe particles (<10 pm) are presented in Figure 2. Preliminary testing suggested much lower reactivity for the Aldrich iron particles. Therefore, a much higher metal to solution ratio (10g/20mL) was used for the reaction. 0.103 mM (15.9 mg/L) of CT was reduced within about 72 hours (>98%). Three major products were identified. CF peaked (53%) around 48 hours. Complete reduction of CF was observed around 100 hours. Concentration of DCM increased rapidly corresponding to the CT reduction and leveled off around 70 hours at a yield of 71%. Methane accumulated with a final yield around 25%. Small amounts of chloromethane also observed after 72 hours and gradually accumulated to about 5% of the total carbon Figure 1. Reactions of nanoscale iron particles with CT. Initial CT concentration was 0.103 mM. Metal to solution ratio was 0.25 g/20 mL
To bet ter compare the reaction rates observed for various iron particles, it is essential to quantify the reactivity per unit metal surface area. The rate of transformation for a chlorinated organic compound in a batch system can be described by the following equation : pseudo-first-order kinetics. Best-fit values of KSA are 5.31X1 O'4 for nanoscale iron and 1.0x1 O'4 for the Aldrich iron, respectively. Several factors may contribute to the difference in reactivity. Laboratory synthesized nanoscale iron surface may have “fresher” metal surface due to less surface oxidation or surface contamination. Mass transfer resistance was also less significant for the nanoscale iron batch system. The metal to solution ratio for the Aldrich iron experiment (10 g/20 mL) was 40 times of that for the nanoscale iron experiment (0.25 g/20 mL). However, the two batch systems had similar mixing intensity (mixed on a rotary shaker at 30 rpm). It was observed that most of the Aldrich iron particles were settled at the bottom of the bottle while Figure 2. Reactions of commercial grade iron particles (Aldrich, <10 pm) with CT. Initial CT concentration was 0.103 mM (15.9 mg/L). Metal to solution ratio was 10 g/20 mL.
Where C is the concentration of organic compound in the aqueous phase (mg/L), KSA is the surface-area-normalized rate coefficient (L/h/m2), as is the specific surface area of metal (m2/g), pm is the mass concentration of metal (g/L), and t is time (h). For a specific system, KSA, as and pm are constants. The above equation therefore represents a
nanoscale iron particles were essentially suspended in solution. Therefore slow transport or diffusion of chlorinated methanes to the settled Aldrich iron surfaces may have caused the slow reaction for the commercial grade iron particles. A major advantage of the nanoscale particles for treatment of CT is the low yield of DCM. Yield of DCM was merely 55% with the nanoscale Fe, compared to 70% with the Aldrich iron. Reductions of CT by the nanoscale iron particles also tend to be more complete with higher yields of methane. For example, yield of methane was 41% for CT reaction with nanoscale iron particles, compared to only 23% with the Aldrich iron. This study demonstrated the potential of the nanoscale metal particles for transformation of CT. Combining the effects of larger surface area (~39 times) and higher surface reactivity (~5.3 times), performance of the nanoscale system versus conventional zero-valent iron barrier is expected to be appreciably (-206 times) higher. However, our knowledge on the underlying reaction mechanisms at the metal-solution interface is still primitive at the best. Important questions need to be investigated include: (1) the exact role of palladium for the dechlorination, (2) optimal amount of surface palladium coverage, (3) effects of naturally occurring oxidants (e.g., oxygen) and reductant (e.g., sulfide) on the long term performance of nanoscale iron particles, (4) cause for the slow DCM dechlorination. Clearly, for this technology to be fully optimized for environmental applications, a better understanding of the fundamental chemical mechanism is essential. REFERENCES 1. Wang, C. B. and Zhang, W. 1997. “Nanoscale Metal Particles for Dechlorination of PCE and PCBs,” Environ. Sci. Technol., 31(7):2154-2156. 2. Wang, C. B. and Zhang, W. 1997. “Catalytic Reduction of Chlorinated hydrocarbons by Pd/Fe, Pt/Fe, and Pd/Zn Bimetals,” 15th Meeting of the North American Catalysis Society, May 18-23, 1997, Chicago. 3. Lien, H and Zhang, W. 1998. “Transformation of Chlorinated Ethylenes in Aqueous Solution Using Nanoscale Bimetallic Particles,” submitted to Journal of Environmental Engineering. 4. Johson, T.L., M.M. Scherer, and P.G. Tratnyek. 1996. “Kinetics of halogenated organic compound degradation by iron metal,” Environ. Sci. Technol., 30(8):2634-2640.
ALAN F. YEN, Ph.D., P.E., DEE AFY, Inc. 504 Harvard Avenue Swarthmore, PA 19081 EDWARD G. HELMIG
In order to characterize the spent blast material and to study the variability among samples, ten actual ABM samples were collected from Atlantic Marine Corporation, and respective metal concentrations were measured using two EPA standard methods i.e. the toxicity characteristic leachate procedure and total
its compressive strength, that is, its behavior under load. To understand the behavior of hardened concrete with blast material and to study the variability in samples from one operation to another ten different samples (moisture content 0- 9.5 % and finenes of modulus 1.23-3.6) were used. Four samples per blast
Sorbent: Dover DGSL
MODELING ASPECTS OF CONTAMINANT TRANSPORT STUDIES USING A GEOTECHNICAL CENTRIFUGE δ"1 St. 8th St. INTRODUCTION
THOMAS F. ZIMMIE Civil Engineering Department Rensselaer Polytechnic Institute
In general, the two commonly used geotechnical centrifuges are the balanced-arm type and the drum type. The balanced-arm centrifuge consists of the centrifuge arm, centrifuge platform that holds the model, balancing counterweight, various fluid and signal connections for data acquisition and mechanism signaling, and the driving mechanism. This centrifuge will typically
,200 g-tons centrifuge at the U.S. CENTRIFUGE MODELING PRINCIPLES Modeling of Models Scaling Relations
Army Corps of Engineers Waterways Experiment Station in Vicksburg, Mississippi. The use of a centrifuge is advantageous for studying phenomena dominated by the earth’s gravity such as saturated groundwater flow. The basic idea is to
corresponding model distance. An example is shown in Figure 1 for a simple flow problem where a 1/N scale centrifuge model experiment was run for one day, using g = 100. As is shown, the prototype seepage velocity is 1/100 that of the model while the corresponding distance is the model. Under these modeling conditions the time for a fluid particle to
represents the transition zone between laminar and turbulent flow for flow g, however, for sands the R<. is 1 at about 50 g for g for a medium to fine sand. 0, mechanical dispersion will CENTRIFUGE MODELING AS A PREDICTIVE TOOL
through porous media . Thus it is important that Rc remain below 1, so that flow remains laminar in both the model and prototype. Under this condition, Darcy’s law is valid. Under normal circumstances, turbulent flow through most soils is rare. It should be noted that Rc in the model and prototype will not be equal (i.e., no similitude of Reynold’s numbers) since the fluid flow velocity in
 which is Radioactive Contaminant Transport g-ton geotechnical centrifuge . The Using the finite was injected into the soil at position GM 2, traveled through the soil, the radioactive count rates were traveled through the soil. The time of 115 minutes does not
basically a small test pad. This type of testing in the laboratory, under 1 g conditions, is unlikely and if attempted, very time consuming. Analytical and numerical solutions are often used for most studies on contaminant transport. Analytical solutions are more common since they are simple to use. However, in general, analytical solutions tend to be highly
The following conclusions can be made pertaining to the centrifuge modeling of contaminant transport processes: 1. Centrifuge model tests have the ability to accelerate the transport processes (by a factor of N2) in flow problems and provide stress levels similar to prototype stresses.
Figure 2 - Results of batch removal of arsenic [As(V)] with Connelly-GPM iron and calcareous sand (initial concentrations 50 to 22,000 μg/L) 500 j ................................................... O ppb As(V)..................... “Šk 400 Flow = 1.0 mL/mìn « 200 HRT = 6.0 minutes І 100 O V · — *
------------1------------l·-^ 0 200 400 600 800 1000 e Volumes Figure 3 - Results of laboratory column arsenic [As(V)] retention experiment. pH 6.3 and I.S.=0.01 M
f non volatile transformation intermediates. Incorporation o f radiolabelled carbon into biomass was not determined. These and similar data were evaluated to determine the maximum rates o f naphthalene mineralization for the three sorbents and four aging periods (Figure 2). Figure 1. Distribution o f 14"C-components during biodegradation o f
A NEW ION-EXCHANGE PROCESS FOR CHROMATE REMOVAL Dongye Zhao, Arup K. SenGuota and Lori Stewart Environmental Engineering Program 13 E. Packer Avenue Lehigh University Bethlehem, PA 18015 INTRODUCTION
Cr(VI) concentration. Understandably, polymeric anion exchangers in fixed-bed column configuration have been widely used both at laboratory and commercial- scale levels for chromate removals from contaminated waters (7-10). Previous
PAA adsorption on aluminum oxide was found to decrease with an increase in pH (Figure 2). The adsorption is ligand-like and is identical with that noted by other investigators [10,11]. The adsorption shows a monotonie decline with increasing pH. A maximum of 90% adsorption was obtained at the lowest pH value. Similar adsorption curves were obtained in the adsorption of humic and fulvic acids on alumina . The ligand exchange refers specifically to direct bond formation between a carboxylate group and a metal ion center of the surface (Al(III) this case). The adsorption trend observed in ligand-like adsorption results from competition for H+ by anions and the mineral surface and is characterized by the mineral pHzpc and the pKa of the sorbate . PAA has an average pKa of 4.5 . The surface sites of an aluminum oxide particle are protonated and positively charged at a pH value below the pHzpc (8.9 for the aluminum oxide used in this study; they are deprotonated and negatively charged at pH values higher than pHzpc. At lower pH PAA is negatively charged and therefore will be adsorbed strongly on the more energetic available positive A120 3 sites. With increasing pH, as the sites become negative, the PAA percentage adsorption decreases. The second step in this study is to analyze the Cd adsorption on A120 3 in the presence of PAA. The presence of a complexing ligand can enhance or decrease the extent of metal adsorption on oxides. The functionality and adsorptive properties of the ligand determine its effect on metal uptake. In the system containing both Cd(II) and PAA, the surface (aluminum oxide) and the ligand (PAA) compete for the metal ions. The differences in the adsorption of Cd(II) on aluminum oxide in the presence of PAA can be explained by
COD is reduced (the COD removal is 49%) for 25,000 mg/1 hydrogen peroxide and 2,000 mg/1 ferrous ion applied. The wastewater pH is reduced to 1.38. Fig. 3 shows variations of the residual hydrogen peroxide in the wastewater treated with hydrogen peroxide and Fenton's reagent. For the oxidation process using hydrogen peroxide only, the hydrogen peroxide applied is decomposed about 40%. Using Fenton's reagent, a higher percentage (80%) of hydrogen peroxide is decomposed. The reaction between hydrogen peroxide and ferrous ion results in the production of hydroxyl radicals (see eq.l) which are effective to oxidize the organic chemicals that are difficult to oxidize with ordinary oxidants. Thus, the decomposition rate of hydrogen peroxide and COD removal rate for Fenton oxidation are higher than for hydrogen peroxide oxidation. ( 1 )
The variation of hydrogen peroxide is an important parameter for the Fenton oxidation process. Fig. 4 shows the COD removal at different concentrations of hydrogen peroxide. More COD is removed at higher hydrogen peroxide concentrations with 44% of COD removed with 25,000 mg/1 hydrogen peroxide applied.
pero xide is approximately 97-99%. Figure 5. Variations of COD in the paper processing wastewater. Experimental condition: the rate of aeration is 31/min. Hydrogen peroxide oxidation : [H20 2] = 10,000 mg/l. Fenton oxidation : [H20 2] = 10,000 mg/l, [Fe2+] = 2,000 mg/l. Figure 6. Variations of pH in the paper processing wastewater .
Figure 7. Variations of the residual hydrogen peroxide in the paper processing wastewater The oxidation process using only hydrogen peroxide also has good COD removal efficiencies. Fig. 8 shows the COD removal observed at different concentrations of hydrogen peroxide. Higher COD removal efficiencies are observed for higher concentrations of hydrogen peroxide applied. When 10,000 mg/1 hydrogen peroxide is used, the COD removal can reach 59%. Figure 8. Effect of hydrogen peroxide in the paper processing wastewater
Re sults of the Kinetic Experiment Table II provides data for the twelve experimental runs. When initial PCP concentration was 100 mg/1, the best removal occurred with an initial hydrogen peroxide concentration of 0.1M to reduce PCP to 0.03 mg/1 in 120 minutes. With an initial PCP concentration of 10 mg/1, 0.01 M of hydrogen peroxide reduced PCP concentration to 0.03 mg/1 in 15 minutes and to below detectable limits in less than 30 minutes. Finally, the best removal for 1 mg/1 of PCP occurred with 0.001 M of hydrogen peroxide reducing PCP concentration to 0.01 mg/1 in 5 minutes and to below detectable limits in less than 15 minutes. Therefore, the optimum concentration of hydrogen peroxide decreased with decreasing initial PCP concentration. Since all experimental runs were described by the first order reaction kinetics, the removal of PCP generally follows first order reaction kinetics. Models The independent variables of the kinetic experiments were time, and
where ki = First order rate constant; kmax = constant to determine maximum height of the curve; h = constant that determines the spread of the curve; and Oopt = concentration of hydrogen peroxide at which ki is maximum.
The equa tion is modified as follows to account for the influence of (3) where a,b,c,d = constants. The model was applied to the data, resulting in 99 .115 % of the variance being explained, an R of 0.996 and a final loss of 0.00437. The determined constants were as follows: a = 1.693, b = 0.688, c = 396.056, d= 0.862, h = 0.301, and kmin = 0.0195. That is, the model is good for the range of PCPj from 1 mg/1 to 100 mg/1 and from 1 M to 0.001 M. Therefore, the model is only good for
the rate constant for benzene oxidation by both hydrogen peroxide and UV light the benzene concentration, μΜ; and the hydrogen peroxide concentration, μΜ. a, b, and c are reaction orders Again, the data was applied to the above model. It was found that this
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Norwich, a small city in Eastern Connecticut located in New London County, is at the source of the Thames River that flows into Long Island Sound passing the U.S. Submarine base in Groton and the Coast Guard Academy in New London along its way. The area contains several other small cities, most of which supply some of their residents with municipal water. Small towns also spread over the 600 sq. miles of county land along with some farms and many residential areas. Gradually over the years a considerable number of well water tests began to show pH's below the desirable range of 6.5 to 9.0. A summary of these pH tests and the increasing acidity trends are shown in Figure 1. for two types of wells, dug and drilled.
Evidence of corrosion in home piping system began to appear in fixture stains and metallic taste. In 1985, the Federal Government did surveys of heavy metals in water supplies, focusing on lead in drinking water and its effect on children. The prevalence of this metal resulted in the banning of lead from solder used in plumbing. In 1988, and in 1989, regulations reduced the safe limit of lead in water from 50^gm/L to 5 μgm/L, with an action level at 15pgm/L in a first draw sample. At the same time the E.P.A. provided guidance documents to schools and day care centers to help reduce children's lead exposure in water supplies. Grants were made available to states to be used for spreading the knowledge about this danger. However, the Federal regulations applied only to public water supplies, which were defined as systems serving 25 or more people. But private wells serving family homes, many containing children, were not included. Children could be assured of regulated water supplies in school, but not at home. Aware that lead was a problem in paint chips, Connecticut required that pediatricians test for traces of lead in children at age two. The Federal act recognized that the primary source of lead in water supplies came through home plumbing systems stemming from the corrosion by the water. This was remedied by requiring public water supplies to prevent corrosion, usually by adding alkalinity to the water. Here again, private wells not included in the educational phase of the program were also not included in the remedies. While these activities with water supplies were occurring, a phenomenon of a different sort was originating in states hundreds of miles away from Southeastern Connecticut. Throughout the Midwest and into West Virginia and Kentucky, electric power generators were erecting tall stacks to disperse sulfur dioxide gasses high into the atmosphere. The gas came mostly from the burning of high sulfur coal mined nearby. The Federal air pollution standards for ambient air were being met in the Midwest by the use of these tall stacks. In the eastern states, restrictions on sulfur content of fuels, mostly petroleum based, were used to meet the ambient air standards. By the early 1980's Connecticut eased its restrictions on sulfur content in these fuels from 0.5% to 1% by weight . Still, during the next five years the sulfur dioxide level actually decreased. However, in 1986, other data collected by the State showed that 32% of the rain storms had an acidic pH of 4.0 or below [ 1 ]. The lowest pH ’ s recorded that year were 3.6. The State also reported that from 1985 to 1996 there had been a further decrease in ambient sulfur dioxide levels [ 2 ].
Un fortunately, the State's acid rain testing program was stopped in 1990. But unpublished data collected by the author in recent years, and reports from New York State strongly suggest that the acidity of rain has not significantly improved. ORIGINS AND EFFECTS OF ACID RAIN
Sulfur dioxide and nitric oxide are the two main factors in the formation of acid rain. When high sulfur coal is burned in Midwestern and other coal producing areas, the combustion gas vented through tall stacks travel with the normal west to east winds to
and soils. They also do not have the problem of acid wells. Therefore, it should be recognized that the acidity of Southeastern Connecticut's ground water is due to a combination of acid rain and the native rock and soil properties. Figure 2 is a picture taken in the Norwich area, at Exit 81 on Interstate 395. This outcrop shows the typical bedrock of the county. Figure 2 Typical Granitic Rock Outcrop in Norwich Conclusions and Recommendations
Federal and state environmental agencies need to recognize that acid rain is contributing to the acidity of both well waters and surface waters. Both are
th at originates in neighboring countries has been recognized as a cause of well water problems. The Swedish government has alerted home owners to the danger, and even subsidized the purchase of home water treatment systems . In New York, where the Adirondack lakes continue to be affected by acid rain. Long Island Lighting Company (LILCO) announced it will no longer sell its pollution rights to Midwestern or Southeastern power companies . Recently many states, including Connecticut, have passed electric deregulation bills. This encourages large scale users to purchase the cheapest power available. Some citizens are concerned this will increase the burning of coal in plants without pollution controls and lead to more acid rain. Perhaps the only recourse for citizens will be to come together and initiate Federal court actions to prevent downwind environmental damage. References
1. "Annual Air Quality Survey for 1986." Department of Environmental Protection, State of Connecticut. 2. "Environmental Quality in Connecticut, 1996 Annual Report." Council of Environmental Quality. 3. Rogers, John. 1985. "Bedrock and Geological Map of
Department of Environmental Engineering and Health Chia Nan College of Pharmacy and Science Tainan 717. Taiwan, R.O.C.
(9) the oxidation rate of organic compounds is fast when large amount of ferrous ions are present because large amount of hydroxyl radicals are produced. However, the Fenton reaction may slow down due to the slow ferrous ion production. In previous studies, the photocatalytic degradation of dichlorvos (DDVP. an insecticide) on glass supported titanium dioxide was investigated. Results indicate that photocatalysis can be an effective process for the degradation of dichlorvos . The mineralization of dichlorvos and the reduction of toxicity were investigated via the photocatalytic reaction . This study uses Fenton’s reagent, ferrous ions/hydrogen peroxide, to oxidize dichlorvos with an attempt to explore the behavior of dichlorvos oxidation and how factors such as pH, [H O ],
The reaction rate in Equ. 6 is much slower than that in Equ. 1. It is derived that ferrous ions are exhausted quickly, but reproduced slowly . Consequently,
reacting with hydrogen peroxide to produce hydroxyl radicals is 53 and the reaction constant of ferric ions reacting with hydrogen peroxide to form ferrous ions is 0.02 . Therefore, it can be derived that the former reaction is far more swiftly than the latter, resulting a higher rate of hydroxyl radical formation in the first stage reaction than that in the second stage reaction. The rate of dichlorvos oxidation in the second stage is slower than that in the first stage, and the second stage will be referred to here as the Fe3+/H stage.
Abandoned Mine Lands, XXV 468 Aeration, XXVIII, 531 Abandoned Mine Lands, XXIX, 553 Aerobic, XXIX, 54, 414, 434 Abandoned Wastes, XXV 305 Aerobic Biodegradation, XXVIII, 15, Abatement Costs, XXVI, 380 109, 326, 537
Phytoremediation, XXIX, 605 D., XXVII, 528 Pilot Scale, XXVIII, 611, 619 J. C., XXVI, 285 W O., XXVII, 274 Pentachlorophenol, XXY 209 Plant Survey, ΧΧΥ 159 Pentachlorophenol, XXVI, 637 Plastics, XXVI, 557 Pentachlorophenol, XXVIII, 655 Platinized Titanium Dioxide, XXIX,