2 The 4 Technology Solutions
If you have comments or suggestions, e-mail me at email@example.com
A Review of Chemical Oxidation Technology
The use of chemical oxidation for the in-situ remediation of soils and groundwater impacted with Constituents of Concern (COCs) is a technology that has seen significant development and application in the past decade with numerous site applications in the last five years. Three chemicals are typically utilized as oxidants:
- Hydrogen peroxide in the form of Fenton's reagent
- Potassium or sodium permanganate
Each has advantages and disadvantages, as does the use of chemical oxidation technology in general. There are numerous recent reviews of field application of the technology (EPA, 1998 and ESTCP, 1999). The purpose of this review is to assess the utilization of in-situ chemical oxidation in general and the use of these three chemical oxidants in particular. In addition these are powerful chemicals and care should be taken with their shipping, storage and use.
Chemical oxidation has had a long history of application in the field of waste water treatment. Hydrogen peroxide, Fenton's reagent, potassium permanganate, ozone and combinations of ozone with other oxidants are all chemical oxidation systems that have found application for the treatment of waste water. One of the key design issue for the application of chemical oxidation to waste water is the total carbon load over and beyond that represented by the COCs that are targeted in that waste stream. The chemical oxidants will react with and be consumed by all constituents in the waste water stream not just the COCs. In the case of the in situ treatment of groundwater there are also consuming reactions associated with the geologic mineral matrix. In many instances this adversely impacts the cost of application to the point of being impractical.
The general value of in-situ chemical oxidation technologies resides in two areas: first the treatment of residual free product and secondly reduction of overall remediation time frames. Specifically in-situ oxidation is likely to be selected for difficult applications that include:
- Low-permeability soils;
- Highly stratified soils;
- Low solubility compounds;
- High concentrations of highly soluble organics (such as ketones, alcohols or MTBE) that would be difficult to treat with conventional surface treatment technology (air stripping or activated carbon).
- Target compounds with low in-situ degradation kinetic constants; and
- Dense, non-aqueous-phase liquid (DNAPL).
From an economic perspective chemical oxidants are usually only practical in limited areas, typically near or in source zones. For the purpose of this review it is important to differentiate types of sources zones:
- Primary source zones refer to areas that have been exposed to free phase NAPL. Discrete NAPL pools, emulsified NAPL, and high levels of adsorbed COCs are characteristic of such zones.
- Secondary source zones refer to areas that have been exposed to high concentrations of dissolved COCs. Elevated levels of adsorbed and dissolved COCs are characteristics of these zones.
Chemical oxidation systems have the potential to offer rapid (from a week to a month) removal of COCs. In addition, oxidation systems can be applied to dissolved plumes that are above the levels for which natural attenuation mechanisms can be expected to work within the time frames desired for site closure. Treatment of large diffuse dissolved COC plumes is possible with chemical oxidants, but often not economically practical.
The most powerful advanced oxidation systems are based on the generation of hydroxyl radicals. The hydroxyl radical is an extremely powerful oxidation agent, second only to Fluorine in power. Fenton's reagent, ozone, and sonification are different means of generating those hydroxyl radicals. Each has advantages and disadvantages that can be exploited given the specific setting of a COC impacted site. Less intense, but still applicable oxidation method, involves the use of potassium or sodium permanganate. Following is a listing of common chemical oxidants, placed in the order of their oxidizing strength:
Relative Oxidation Power (C12 = 1.0)
Hydroxyl Radical (OH) 2.06
Atomic Oxygen (singlet) 1.78
Hydrogen Peroxide 1.31
Perhydroxyl Radical (OOH) 1.25
Potassium Permanganate 1.24
Chlorine Dioxide 1.15
In addition, the geology must be considered when applying a chemical oxidant in-situ. While the first concern is having the liquid reagents reach all of the COC in a heterogeneous situation, the soils themselves can have a significant effect on the reactions. This technology relies on water as a carrier to delivery the chemical oxidant to COCs. The macro and micro geologic conditions do not always allow for complete contact between the chemical oxidants and the COCs present in the vadose or saturated zone.
It is important to understand that COC oxidation occurs only in the aqueous phase, involving dissolved species of both the COC and the oxidant. The solubility of a targeted COC ultimately controls the rate of possible oxidation. There is interface mass transfer and then chemical oxidation. The rate of mass removal is limited by the kinetics of the COC dissolution process rather than those of the chemical oxidation reaction.
In all cases where the final product of the oxidation product is carbon dioxide there is potential for short term plugging of the aquifer pore space with carbon dioxide gas bubbles. In the case of Fenton's reagent, oxygen gas will also be generated and contribute to this possible problem.
Lastly, with the exception of some permanganate applications, this technology relies on the total displacement of native groundwater from the treatment zone and replacement with the reagent laden water. In the case of permanganate, the active life is sufficient that at sites with moderate to high groundwater velocities (greater than 30 feet per year) reagent injection can take place with subsequent reliance on advection and dispersion to mix the oxidation solutions with native groundwater during transport.
Chemical Oxidation with Fenton's Reagent
Hydrogen peroxide alone has little effectiveness for the oxidation of chlorinated solvents.
The Fenton's reaction was first described in 1894, during the 1930's the reaction mechanisms were fully defined, over the last 15 years commercial reactors have been available for waste water treatment, and in 90's applications for in-situ groundwater treatment have been developed. Fenton's reagent generates hydroxyl radicals through the reaction of ferrous iron and hydrogen peroxide:
Fe+2 + H2O2 ==> .OH + Fe+3 + OH-
The process is self replicating since the reaction of ferric iron with hydrogen peroxide to generate the perhydroxyl radical also occurs:
Fe+3 + H2O2 ==> Fe+2 +.OOH + H+
The perhydroxyl radical is a weaker oxidizer (between hydrogen peroxide and permanganate). But, more importantly the process generates ferrous ions that in turn stimulate further reaction with hydrogen peroxide to produce more hydroxyl radials. The hydroxyl radical can react with almost any hydrocarbon to produce carbon dioxide as a final product (and chlorides if a chlorinated hydrocarbon is treated). The key application issues are delivery of the reagents in-situ and pH control. Optimum pH for the reaction, in the range of 3.0 to 4.5, is driven by the iron chemistry It is possible to use iron chelates to get the desired reaction at higher pH's (Sun and Pignatello, 1992). Improper control of the reaction will only generate oxygen and water, not high intensity oxidation. The technology is ideal for application in source zone areas where there are high concentrations of adsorbed or even interstitial free product. Cost is mostly dependent upon the geochemistry of the geologic matrix and control of the iron chemistry. Soil buffering capacity may limit the ability to obtain optimal subsurface pH conditions.
Most of the original utilization of Fenton's reagent was in relatively low concentrations in wastewater applications. The main criteria for wastewater was to have a sufficient ratio of hydrogen peroxide to the organic chemical that was to be destroyed. The chemical oxidation process is less efficient in the presence of soil. Due to the presence of competing organics and mineral surfaces (Miller and Valentine, 1995) that are reactive to hydrogen peroxide, and the less than optimal environmental conditions, dosage requirements for in-situ applications may increase 10 to 100 fold or more to accomplish the desired oxidation of COCs. With regards to the rate of reaction and the degree of completion of the reaction, high concentrations of hydrogen peroxide (10 to 25%) are preferred. On the other hand, excess consumption of hydrogen peroxide by reactions with the geologic matrix can be controlled by using applications (possibly multiple) at concentrations of 5% or less.
Almost all soils have native iron bearing minerals as part of their geochemical makeup. When soils contain chemically available iron, supplemental iron salts may not be required with the hydrogen peroxide injection solution (Teel et al, 2001). However, problems may occur when too many iron minerals are present in the soil, if the iron is in a mineralized form not readily available for dissolution under low pH conditions, the natural catalytic activity of the mineralized iron decomposes the hydrogen peroxide to oxygen and water and does not create the hydroxyl radical. A second problem with soil mineralogy comes from soils rich in carbonate minerals or with high alkalinity. These soils produce groundwater with a high pH. This can required unacceptable amounts of acid to buffer the pH down to the required range. Carbonate minerals respond to low pH so rapidly that the geologic matrix itself will consume acid. Lastly, the carbonate ion preferentially scavenges hydroxyl radicals before they have a chance to react with the COCs.
The heat generated by exothermic dissociation of hydrogen peroxide promotes contaminant volatilization. Rapid gas generation creates turbulence that further enhances the contact of the oxidant with the targeted COC impacted zones. Concentrations as low as 11 percent can cause ground water to boil. The rate of hydrogen peroxide decomposition doubles with every 10 degree C rise in temperature so the energy release process, once initiated, rapidly accelerates. A pound of hydrogen peroxide can release 1,200 BTUs of heat energy (1)(2) and up to six cubic feet of oxygen gas. This creates volatilization and transport due to the created pressure gradients. Which when exploited in a controlled fashion can have beneficial effects regarding contact of the oxidant with the COC entrained in the geologic matrix.
For well over a decade the cost of 50% hydrogen peroxide has been $3.50 a gallon.
With the permanganate ion the initial oxidation reactions of COCs are independent of pH. For example, The oxidation rate of chlorinated alkenes by potassium permanganate is not effected by pH over a range of 3.0 to 11.0. However, pH does impact the type of intermediate products produced, at pH 4 formic acid is dominant, at pH 6 to 8 oxalic and glyoxylic acids dominate (Yan and Schwartz, 2000). In both cases the organic acids are subsequently oxidized to carbon dioxide. The conversion of the intermediate oxidation products to carbon dioxide does occur more rapidly under acidic pH conditions. For the initial permanganate oxidation reactions, they can be generally characterized (Wiberg and Freeman, 2000) as follows:
- Acid catalyzed reactions at pH < 5.0
- Uncatalyzed reactions at pH 5-9
- Base catalyzed reactions at pH >10
Potassium permanganate has an affinity for organic compounds containing carbon-carbon double bonds, aldehydes groups and hydroxyl groups. The permanganate ion borrows electron density from the pi bonds in chlorinated alkenes. This creates a bridged oxygen compound known as the cyclic hypomagnate ester (Yan and Schwartz, 2000). The intermediate ester is unstable and further reacts by a number of mechanisms including hydroxylation, hydrolysis or cleavage. Under normal pH and temperature conditions the primary oxidation reactions involves spontaneous cleavage of the carbon-carbon bond. Once the double bond is broken, the highly unstable carbonyl groups are immediately converted to carbon dioxide through either hydrolysis or further oxidation by the permanganate ion.
Manganese dioxide (MnO2) is the end product of the reduction of the permanganate oxyanion under neutral to basic pH conditions by the following reaction:
MnO4- + 2H2O + 3e- ==> MnO2 + 4OH-
The MnO2 is insoluble and forms colloids with a typical diameter near 1 micron (West et al, 1998). Plugging of the effective pore space can occur due to the precipitation of MnO2. The manganese dioxide colloids can also cause problems with surface treatment equipment used as part of oxidation programs that rely on the creation of active circulation cells. The formation of insoluble manganese dioxide reaction products can also form a coating and seal around NAPL mass, in effect blocking further dissolution and migration of the COC into the adjacent reaction areas (Urynowicz and Siegrist, 2000). This can be a significant problem.
At low pH conditions the reduction of the permanganate ion proceeds to Mn(II) by the following reaction:
MnO4- + 8H+ + 5e- ==> Mn+2 + 4H2O
The Mn (II) cation is soluble in water at concentrations above the regulatory limit (50 ppb) when chloride or sulfate counter anions are present.
Examples of complete reactions of permanganate with CVOCs are as follow:
- PCE C2Cl4 + 2MnO4- ==> 4Cl- + 2CO2 + 2 MnO2(s)
- TCE C2HCl3 + 2MnO4- ==> 3Cl- + 2CO2 + H+ + 2MnO2(s)
- DCE C2H2Cl2 + 2MnO4- ==> 2Cl- + 2CO2 + 2H+ + 2MnO2(s)
The stochiometry ratios for complete oxidation reactions of permanganate as pounds of the KMnO4 salt required to oxidize 1 pound of CVOC is:
- PCE 1.3 to 1;
- TCE 2.4 to 1;
- DCE 4.4 to 1; and
- VC 8.5 to 1.
(Note the dramatic change in consumption rate is due to the fact that this analysis is on a pound per pound basis, not a molar basis. A pound of vinyl chloride has almost 7 times more moles than a pound of PCE due to the loss of 3 heavy chlorine from ethene backbone).
In contrast to Fenton's reagent, the presence of carbonate minerals in the geologic matrix has generally positive effects on permanganate oxidation (Nelson et al, 2001):
Mn reduction stops at Mn(IV) (insoluble) rather than proceeding to potentially soluble Mn (II).
Trace minerals are coprecipitated and immobilized with the manganese oxide.
Consumption of permanganate by reaction with reduced mineral such a magnetite is minimized because the rates of these reactions increase with decreasing pH.
Oxidant demand from the matrix can be attributed to natural organic matter, reduced metals, carbonates and sulfides. Permanganate demand rates can vary from a few grams of permanganate per kg of soil (clean sand with dissolved COC) to as much as several hundred grams of permanganate per kg of soil (organic clays with 6% organic carbon content and NAPL).
Immobile reduced metals such as chromium, uranium, vanadium, selenium and molybdenum (all of which are soluble as oxidized oxyanionic complexes) can be oxidized and mobilized by permanganate oxidation.
When the permanganate concentration is 10 times the concentration of the CVOCs the half life of TCE and DCE ranges from 24 seconds to 18 minutes, the half life of PCE is 257 minutes. Typical dosage range (by weight) is 1.5% to 5% permanganate in the targeted treatment zone.
The potassium salt has a solubility limit of 65 grams per liter (6.5%), the sodium salt is significantly more soluble. However, the sodium salt is significantly more reactive. The reactivity difference between the sodium and potassium salt is not significant with regards to COC treatment. It is significant because the increased reactivity of the sodium salt presents greater safety challenges during its shipping, storage and application.
Potassium permanganate costs $4.75 per pound.
Ozone has been used for the treatment of drinking water since 1903. Its oxidizing power and innocuous decay products have made it ideal for that application. Ozone can generate hydroxyl radicals via catalytic decomposition of water, the oxidative power of an ozone generated hydroxyl radical is the same as those generated by Fenton's reagent. It also has specific activity towards alkenes (including chlorinated alkenes) through the attack of the double bond carbon in a fashion similar to permanganate oxidation. Ozone bridges the carbon-carbon double bond to form unstable ozonide intermediates, which decompose into smaller oxidation species until carbon dioxide and water, or stable refractory compounds, such as acetic acid or oxalic acid, are formed. Ozone does not react well with chlorinated alkane's, particularly chlorinated methane's. Although the fate and reaction mechanisms of ozone in porous geologic media is not completely understood, it is likely that hydroxyl radical production from ozone occurs through catalytic reactions with natural inorganic (probably iron oxides) and organic material.
Ozone is fairly stable in dry air and has a half-life of several hours in low concentration. In water, ozone half-life is several minutes. Because ozone is very reactive in an aqueous environment, ozone can oxidize material between 10 to 1000 times faster (Hoishe and Bader, 1983) than most oxidants used in water treatment. Because ozone has such a short half-life, it cannot be compressed and stored. Instead, it must be generated on-site and used immediately. Electrical generation is the only practical and safe method for large-scale applications of ozone. In practice, ozone concentrations of 1-2% using air, and 3-8% using oxygen can be obtained by corona discharge generators. At a production concentration of 1% by weight in air 650 ft3 of air is required to produce 1 pound of ozone. Because the speed of ozone destruction by decomposition is proportional to the ozone concentration, producing higher ozone concentrations with corona discharge is not feasible.
The decomposition rate of ozone has been shown to be 25 times more rapid in geologic media than non-reactive media such as glass beads. Although, the presence of carbonate ions in the groundwater can provide stabilization and double the half life of dissolved ozone at a pH of 8.0 (Staehelin and Holgne, 1982).
Other species created from ozone include:
At atmospheric pressure the solubility of ozone is 3.4 mg/L for 1% and 7 mg/L for 2%, note that solubility is somewhat dependent upon of the concentration of the ozone in the gas phase. With decreasing temperature solubility increases. In addition, increasing hydrostatic pressure increases ozone solubility, with a 1% system for example:
At 18 psi ozone solubility is 4.2 mg/L;
At 25 psi it is 6.3 mg/L; and
At 32 psi, 8 mg/L.
As the TDS of the treated water increases, the solubility decreases. For example ozone solubility is reduced 30% in sea water (Kosak-Channing and Heiz, 1983).
In most applications 1.5 to 3 pounds of ozone is required for each pound of COC to be treated. That consumption rate is based purely on COC degradation requirements, spontaneous decay and ancillary reactions with other groundwater of mineral matrix constituents create an ozone loading requirement over and above those base line numbers.
The cost for the generation of ozone includes the cost of the equipment:
1 pound per day unit - $7,500;
10 pound per day unit - $32,500;
40% of cost is for air pretreatment.
The power requirements are in the the range of 8 to 12.6 Kw-hr per lb of ozone produced.
Practicality of Application
Stoichiometric Cost Comparison
The weight ratio in terms of pounds of oxidant per pound of TCE fully oxidized are 0.8 lb/lb for Fenton's reagent and 2.4 lb/lb for potassium permanganate. The cost for Fenton's regent including hydrogen peroxide, and amendments for pH control and ferrous iron addition is typically $1.10 per pound and potassium permanganate is $4.75 per pound. So total oxidant cost per pound of TCE destroyed would be $0.88 for Fenton's reagent and $11.40 for potassium permanganate.
Calculation of a similar cost for ozone oxidation is more complicated, involving the capital costs of the equipment, the operational life of the system, electrical efficiency and electrical costs. These are assumptions for those key parameters:
A 10 pound per day ozone unit is purchased for $32,000 for project life of 3 years. This translates to a capital cost of $2.90 per pound of ozone.
It requires 10 Kw-hr at $ 0.10 per Kw-hr (or $1.00) for the electrical cost to produce one pound of ozone.
It requires 2.25 pounds of ozone to oxidize one pound of COC, translating into a total cost of $8.75 for the ozone to oxidize one pound of COC.
Actual Application Costs
The total costs for the use of these oxidants is not accurately reflected by just the pound per pound cost required for the stoichiometric oxidation of the COCs. In reality reactions with native carbon in the geologic matrix, reactions with minerals in the geologic matrix (Barcelona and Holm, 1991), and reactions with native dissolved constituents in the groundwater can increase the consumption rate of the oxidation reagents by orders of magnitude. Fenton's regent and ozone are particularly susceptible to these effects, potassium permanganate less so.
In practical application these chemical oxidants are not applied based on the amount of COC present (with the exception of instances when there are significant volumes of NAPL present), rather they are applied to achieve a specific reagent concentration in the treatment zone. Assuming the desired reagent concentration is 5% for Fenton's reagent and permanganate (this argument is not valid for ozone as explained below), following are the per cubic yard cost for one treatment with the respective chemical oxidant:
In practice it is likely that 2 to 5 applications of Fenton's reagent may be required to achieve the remediation goals during a treatment campaign, raising the reagent costs to the range of $ 42 to $ 105. In many instances one application of potassium permanganate will be adequate for treatment, but in cases where NAPL is present 2 or 3 applications may be required.
The hydrogen peroxide in Fenton's reagent can have a reaction life (under practical conditions) that ranges from ¼ hour to several hours. The permanganate ion can remain active in the saturated zone for months or perhaps even a year or more (Nelson et al, 2001). Ozone will decompose in water within minutes. This has significant impact for in-situ applications:
Fenton's reagent can be directly injected into an aquifer, but spacing of the injection points will typically range from 1 to 7 yards. All contact with the COC occurs during advection caused by the injection process, and by mixing caused by heating and gas generation. Groundwater in the treated areas is displaced.
Permanganate salts can be injected in-situ with the anticipation of further mixing of the reagents with the COC impacted groundwater due to natural flow conditions. Spacing of injection wells can be predicated on the impact of natural groundwater flow over a 6 month to 1 year period. Push-pull circulation systems can also be utilized to decrease the spacing of injection wells, but accommodations must be made for the removal of colloidal manganese dioxide from recovered groundwater. The reaction rate of permanganate also allows it to be applied in conditions where native groundwater flow can provide further advection for transport and dispersion for mixing. Complete displacement of native groundwater is not required.
Ozone is so reactive in water that it can not be practically injected in-situ. It can be used as a reactant in the bore of circulation well or as part of a sparge system. The radius of influence of such systems are limited. As a result the actual application of ozone as an in-situ chemical oxidation system has been much more limited than the use of Fenton's reagent or potassium permanganate.
Conclusion and Cautions
A last word of caution. These are powerful chemicals. During shipment, storage, and application of chemical oxidants great care must be taken to maintain safe site conditions with both regards to personnel and property.
High strength hydrogen peroxide (greater than 5%) can cause chemical burns to the skin and eyes, lower concentrations will be irritating.
In all cases, handling requires protective clothing: Face shield, gloves, hard hat, rubber or PVC boots, and a rubber or PVC suit.
A shower and eye wash should be available.
Decomposing hydrogen peroxide rapidly generates heat, gas, and pressure.
Storage vessels and piping runs that have valve traps should have adequate ventilation and pressure relief systems.
Containers and piping should be free of all contaminants.
Contaminants include: copper, brass, zinc, mild steel, synthetic rubbers, polypropylene, and organic compounds (especially liquid organics).
Acceptable materials for storage include: aluminum, stainless steel, glass, ceramics, Teflon, polyethylene, Viton, and (for temporary storage) PVC.
All metal components must be properly passivated before use.
Check valves are required to prevent back flow into hydrogen peroxide piping or tanks.
Hydrogen peroxide pumps should only be constructed of stainless steel or Teflon.
Hydrogen peroxide can generate enough heat and oxygen to ignite combustible materials. Do not store hydrogen peroxide on wooden pallets or decks.
Hydrogen peroxide in the presence of hydrocarbon vapors can cause vapor phase explosions.
Every 10oC increase in temperature doubles the reaction rate of hydrogen peroxide, it must be stored away from heat sources (boilers, steam lines, etc) and should never be stored in insulated tanks.
Shipment of solutions with a concentration greater than 8% must by as DOT hazardous.
Hydrogen peroxide is sold as a 50% solution. This is an extremely reactive material. Hydrogen peroxide vendors can readily deliver more dilute solutions. When possible it is preferable to have the hydrogen peroxide delivered at the concentration at which it will be used for in situ chemical oxidization, rather than delivered at 50% and then diluted on site.
Permanganates present a fire an explosion risk when placed in contact with organic materials.
Containers or permanganate salts exposed to liquid hydrocarbons may explode.
Other materials that exhibit high reactivity include: metallic powders, elemental sulfur, hydrochloric acid, hydrazine, hydrogen peroxide, and metal hydrides
Permanganate salts are irritating to the eyes, skin, respiratory system and digestive tract from acute exposure.
Chronic exposure to the skin can cause defatting and dermatitis. Chronic ingestion can cause central nervous system and kidney damage.
Employee protective equipment should include: dust and mist respirator (and or area ventilation), gloves, and safety goggles.
Shipping is regulated under the Transport of Dangerous Goods Act.
TDG Classification: Class 5.1 (9.2) U.N. 1490; Packing group II
WHMIS Classification: C, E
On the DSL list
Ozone can not be shipped, it must be generated at the point of use.
Ozone can only be generated at limited concentrations, the chief hazard is exposure to ozone inhalation.
Ozone can be detected by smell at a concentration of 0.01 to 0.005 ppmv.
However, repeated exposure increases the detection level.
For eight hours the exposure limit is 0.1 ppmv.
For 10 minutes the exposure limit is 0.3 ppmv.
Ozone equipment and process lines should be in well ventilated areas.
Exposure to aqueous ozone solutions can cause skin or eye irritation.
Table 1 - Characteristics of Hydrogen Peroxide
- Wgt. % O2
- Wgt. % O2
- Mole Fraction
- Mole Fraction
- Molecular Wgt.
- Molecular Wgt.
- Apparent pH
- Apparent pH
- Ht. Dec. cal/gm
- Ht. Dec. cal/gm
TABLE 2 - Ferrous Iron (Fe II) Containing Minerals
Table 3 - 2nd Order Decay Constants (M-1s-1) -
H2O2/Fe" Ferrous Iron pH 2 0.044
" Ferric Iron pH 2 0.024 - 0.054
" Ferric Iron pH 2.8 0.009 - 0.09
" Ferric Iron pH 2 - 4 0.002 - 0.01
" Fe2O3/Al2O3 pH 12 0.037
" Fe2O3/Al2O3 pH 9 0.013 - 0.031
" Goethite pH 7 0.032
" Goethite pH 7.7 0.0016
" Goethite pH 5 10 0.019 - 0.067
" Ferrihydrite pH 7.7 0.023
Barcelona, Michael J. and Holm, Thomas R., 1991. Oxidation-Reduction Capacities of Aquifer Solids, Environ. Sci. Technol., Vol. 25, No. 9, pp. 1565-1572, 1991.
EPA, 1998. Field Application of In Situ Remediation Technologies: Chemical Oxidation, EPA 542-R-98-008, 31 pp., September 1998, http://www.epa.gov/swertio1/download/remed/chemox.pdf
ESTCP, 1999. Technology Status Review In Situ Oxidation, November 1999, 42 pp., http://www.estcp.org/documents/techdocs/ISO_Report.pdf
Hongine, J. and Bader, H., 1983. Rate Constants of Reactions of Ozone with Organic and Inorganic Compounds in Water 1: Non-Dissociating Organic Compounds, Water Res., Vol. 17, pp. 173-183, 1983.
Kosak-Channing, Lynn F. and Heiz, George R., 1983. Solubility of Ozone in Aqueous solutions of 0-0.6 M Ionic Strength at 5-30 oC, Environ. Sci. Technol., Vol. 17, No. 3, pp. 145-149, 1983.
Miller, Christopher M. and Valentine, Richard L., 1995. Hydrogen Peroxide Decomposition and Quinoline Degradation in the Presence of Auqifer Material, Wat. Res., Vol. 29, No. 10, pp. 2353-2359, 1995.
Nelson, Matthew D., Parker, Beth L., Al, Toma A., Cherry, John A., and Loomer Diana, 2001. Geochemical Reactions Resulting from In Situ Oxidation of PCE-DNAPL by KmnO4 in a Sandy Aquifer, Environ. Sci. Technol., Vol. 35, No. 6, pp. 1266-1275, 2001.
Staehelin, Johanes and Holgne, Jurg, 1982. Decomposition of Ozone in Water: Rate of Initiation by Hydroxide Ions and Hydrogen Peroxide, Environ. Sci. Technol., Vol 16, No. 10, pp. 675-681, 1982.
Sun, Yunfu and Pignatello, Joseph J., 1992. Iron (III) Chelates as Catalysts for Chemical Degradation of Pesticide Wastes by Hydrogen Peroxide, Abstract, Division of Env. Chem., AMS, San Francisco, CA, April 1992, pp. 480-481.
Teel, Amy L., Wargberg, Christopher R., Atkinson, David A., and Watts, Richard J., 2001. Comparison of Mineral and Soluble Iron Fenton's Catalysts for the Treatment of Trichloroethylene, Wat. Res., Vol. 35, No. 4, pp. 977-984, 2001.
Urynowicz, Michael A. and Siegrist, Robert L., 2000. Chemical Degradation of TCE DNAPL by Permanganate, in: Chemical Oxidation and Reactive Barriers: Remediation of Chlorinated and Recalcitrant Compounds, The Second International Conference on Remediation of Chlorinated and Recalcitrant Compounds, Monterey, CA, May 22-25, 2000, Vol. C2-6, pp. 75-82.
West, O.R., Cline, S.R., Holden, W.L., Gardner, F.G., Schlosser, B.M., Thate, J.E., Pickering, D.A., and Houk, T.C., 1997. A Full-Scale Demonstration of In Situ Chemical Oxidation Through Recirculation at the X-701B Site Field Operations and TCE Degradation, ORNL/TM-13556, 101 pp.
Wilberg, Kenneth B. and Freeman, Fillmore, 1999. Kinetics of the Base-Catalyzed Permanganate Oxidation of Benzaldehyde, J. Org. Chem, Vol. 65, pp. 573-576, 2000.
Yan, Y. Eugene and Schwartz, Franklin W., 2000. Kinetics and Mechanisms for TCE Oxidation by Permanganate, Environ. Sci. Technol., Vol. 34, No. 12, pp. 2535-2541, 2000.
Return to the Environmental Technology Index Page
Copyright 2002 David B. Vance
All Rights Reserved
If you have comments or suggestions, e-mail me at firstname.lastname@example.org