Ind. Eng. Chem. Res. 1997, 36, 1559-1566
1559
Hydrothermal Oxidation of Organic Wastes Using Ammonium Nitrate Petra I. Proesmans, Li Luan, and Steven J. Buelow* Los Alamos National Laboratory, Chemical Science and Technology Division, CST-6, Mailstop J567, Los Alamos, New Mexico 87545
The use of ammonium nitrate as an oxidizing agent in hydrothermal oxidation of organic compounds was investigated. The oxidation of model compounds, methanol, acetic acid, and phenol, was studied at 500 °C and 345 bar. High organic, ammonia, and nitrate removal was achieved at stoichiometric concentrations. The oxidation of ammonia by nitrate was much faster than the oxidation of either methanol or acetic acid and only slightly faster than phenol. Nitrogen products included N2, N2O, and NO2- as well as toxic NO and trace amounts of NO2. Carbon products were CO2, HCO3-, CO32-, and CO. The co-oxidant system with hydrogen peroxide and ammonium nitrate was studied to eliminate the NOx production. Stoichiometric concentrations of hydrogen peroxide to the carbon concentrations resulted in undetectable NOx levels. Introduction This research was conducted for the U.S. Army Corps of Engineers Research Laboratory. Ammonium nitrate is being studied as an alternative for ammonium perchlorate (AP) as an oxidizing agent in the Department of Defense 1.1 and 1.3 rocket propellants. Because ammonium nitrate suffers from phase instability, stabilizers like potassium dinitramide are added. As the ammonium nitrate technology develops, ammonium perchlorate based propellants may be replaced by those using potassium dinitramide stabilized ammonium nitrate (KDN-PSAN). This increased use of ammonium nitrate or KDN-PSAN will ultimately create a need for environmentally responsible processes to reuse ammonium nitrate extracted from demilitarized rocket motors. One potential application of recovered ammonium nitrate is as an oxidizing agent in the hydrothermal oxidation (HTO) of organic wastes. HTO of organic wastes converts hazardous organic materials to innocuous products, such as carbon dioxide and water, by reacting them with an oxidizer in hightemperature, high-pressure water, typically above water’s critical point (Tc ) 374 °C, Pc ) 221 atm). Under these conditions, both gases and organics are soluble, allowing rapid mixing and fast and complete reactions. HTO converts organics to simple products in just seconds. High destruction efficiencies (>99.99%) have been achieved for a broad range of chemical and metabolic wastes, including difficult examples such as chlorinated and aromatic hydrocarbons (Tester et al., 1992). Many nitrogen-containing organic compounds have been hydrothermally oxidized with high destruction and removal efficiencies, converting organic nitrogen to N2 and trace N2O (Harradine et al., 1993). Ammonia and nitrate are primary constituents or reaction products in several waste streams targeted for treatment by HTO, and there have been several studies of ammonia and nitrate reactions under hydrothermal conditions. These investigations indicate that ammonia reacts slowly with oxygen (Webley et al., 1991) but can react rapidly with nitrate (Dell’Orco, 1994; Bowman and Fulton, 1995; Cox et al., 1992). Nitrate can also react rapidly with organic compounds, converting them to carbon dioxide and water (Cox et al., 1992; Foy et al., 1993). In addition, recent experiments at Los Alamos * Author to whom correspondence is addressed. S0888-5885(96)00171-6 CCC: $14.00
National Laboratory (LANL) have demonstrated that the use of nitrate along with hydrogen peroxide has significant advantages over hydrogen peroxide alone for some mixed ammonia/organic wastes (Luan et al., 1995). Complete destruction of both ammonia and organic wastes was observed at relatively low temperatures ( 0.5 s-1) that the kinetics of the reactions could not be elucidated with the reactor systems used for this study. In the same set of experiments, significant corrosion was observed at the lowest temperatures. Metal analysis showed that chromium and nickel concentrations in the effluent increased by more than an order of magnitude for experiments at 400 °C as compared to those at 500 °C. Therefore, further experiments were conducted at nominal temperatures and pressures of 500 °C and 345 bar, respectively. The methanol concentrations ranged from 0.2 to 1.8 mol/L, acetic acid concentrations from 0.15 to 0.56 mol/L, and phenol concentrations from 0.03 to 0.15 mol/L. Residence times were 30-40 s. The experimental conditions are given in Table 2. Figures 3-5 show the concentrations of organic carbon and ammonia in the effluent as a function of the
Figure 3. Ammonium nitrate (1 mol/L) reaction with varying methanol concentrations (error bars fall within marker).
Figure 4. Ammonium nitrate (1 mol/L) reaction with varying acetic acid concentrations (error bars fall within marker).
carbon content and ammonia concentrations in the feed solution. Both influent and effluent concentrations are expressed as “reducing equivalent”, calculated as the concentration in mol/L multiplied by the chemical oxygen demand (COD). The COD for ammonia is estimated to be 1.5, assuming complete oxidation to N2 (NH3 + 1.5O f 0.5N2 + 1.5H2O). The COD for the organic compounds is also calculated by assuming complete oxidation to CO2 and H2O, yielding a COD of 3 for methanol, 2 for acetic acid, and 2.33 for phenol.
1562 Ind. Eng. Chem. Res., Vol. 36, No. 5, 1997
Figure 5. Ammonium nitrate (1 mol/L) reaction with varying phenol concentrations (error bars fall within marker).
The graphs indicate that, for a 1 mol/L ammonium nitrate solution, the nitrate concentration becomes insufficient for complete oxidation of both ammonia and methanol at a reducing equivalent concentration of ca. 2.5 mol/L (corresponding to carbon concentrations of 0.5 mol/L as methanol, 0.75 mol/L as acetic acid, and 0.64 mol/L as phenol). This agrees with the concentration predicted from the 2.5 mol/L oxidizing equivalents of 1 mol/L HNO3 (HNO3 f 0.5N2 + 0.5H2O + 2.5O). At higher carbon concentrations, the reducing equivalents of carbon in the effluent increased linearly. For methanol and acetic acid (Figures 3 and 4), the reducing equivalent as carbon increases with a slope close to 1 and 0.7, respectively, while the increase of ammonia in the effluent is negligible. When the organic reducing level is greater than 1.0 mol/L, the organic must compete with ammonia for the nitrate oxygen. Hydrolysis experiments showed that hydrolysis and/or pyrolysis reactions were not important reaction pathways for methanol, acetic acid, and phenol destruction at the temperatures and residence times examined in this study. However, Figure 5 shows that phenol reacts with nitrate nearly as rapidly as ammonia. The limited data set shows that, at excess reducing levels, the concentration of ammonia reducing equivalents in the effluent initially increased at about the same rate as the phenol reducing equivalent. Thereafter, ammonia reaches an equilibrium concentration of near 0.25 mol/L reducing equivalent (0.17 mol/L of NH3). If the ammonia and phenol oxidations would occur equally fast and ammonia and phenol oxidation kinetics would be similar, the ammonia concentration would reach an equilibrium concentration of about 0.2 mol/L. As may be concluded from the figures, phenol appears to be more reactive with nitrate than either methanol or acetic acid. Similar findings were reported in the literature for the oxidation of organic compounds by oxygen. Boock and Klein (1993) compared the relative rates of oxidation reactions (with oxygen) between ethanol, 1-propanol (primary alcohols), and 2-propanol (a secondary alcohol) and acetic acid and methanol. They concluded that ethanol and 1-propanol oxidized most rapidly, followed by the secondary alcohols (2propanol). Methanol and acetic acid oxidized most slowly. The oxidation of acetic acid and 2,4-dichlorophenol by hydrogen peroxide and oxygen was compared by Lee and Gloyna (1990). Much higher conversions were obtained for the aromatic compound compared to acetic acid, with both oxygen and hydrogen peroxide as oxidizer. Carbon products in the liquid phase were found to be mainly the unreacted compounds (methanol, acetic acid,
Figure 6. CO and H2 production in the reaction of methanol with ammonium nitrate.
Figure 7. CO and H2 production in the reaction of acetic acid with ammonium nitrate.
Figure 8. CO and H2 production in the reaction of phenol with ammonium nitrate.
or phenol, for the experiments at substoichiometric nitrate concentrations), trace amounts of formate (for the acetic acid experiments), carbonate, and bicarbonate. Gaseous carbon products were CO2, some CH4, and CO. Figures 6-8 show the correlation between the hydrogen and carbon monoxide production and the carbon concentration in the feed. For both methanol and acetic acid oxidation, the hydrogen and carbon monoxide increased simultaneously for a reducing equivalent concentration in the feed of 2.5 mol/L or more. In the case of phenol, hydrogen production was detected for a feed concentration of higher than 2.5 mol/L, but it slowly diminished for increasing inlet concentrations. Nitrogen products in the liquid effluent depended upon the relative concentrations of all components in the initial feeds. The products included small amounts of nitrate, nitrite, and/or ammonia. Gaseous nitrogen products included N2O, N2, NO, and trace amounts of NO2. Figures 9-11 show the gaseous nitrogen products as a function of the organic carbon content in the feed solution. Overall, the N2 production increased and the
Ind. Eng. Chem. Res., Vol. 36, No. 5, 1997 1563
Figure 9. Gaseous nitrogen products in the reaction of methanol with ammonium nitrate.
Figure 10. Gaseous nitrogen products in the reaction of acetic acid with ammonium nitrate.
Table 3. The ammonium nitrate concentration was 1 mol/L, the H2O2 concentration was varied from 0.3 and 1.2 mol/L, and the organic carbon concentration was between 0.3 and 0.4 mol/L for all experiments. Ammonia and organic removal was greater than 99.9% in all experiments. The nitrogen products found in the liquid effluent included nitrate and nitrite. The concentrations of these species increased with increasing H2O2 concentration. Gaseous nitrogen products were N2, N2O, NO, NO2, and HNO2. The ratio of N2 to N2O gaseous effluent was between 2 and 9 for all experiments. As shown in Figure 12, NO production decreased with increasing H2O2 concentration. The methanol concentration was 0.4 M, with an oxygen demand of 1.2 M, while acetic acid had an oxygen demand of 0.75 M. Only for stoichiometric H2O2 concentrations was the NO concentration negligible. The reactions of phenol produced measurably less NO than the reactions of methanol and acetic acid. NO2 was not quantitatively analyzed, but the ratio of NO2 to NO observed decreased with decreasing initial hydrogen peroxide concentrations (NO + 0.5O2 S NO2 reaction shifted to the left at lower O2 concentrations). HNO2 was observed in the FTIR analysis for the acetic acid and phenol oxidation experiments at the lowest H2O2 concentrations. None of the reaction produced H2, and only the reactions of methanol produced CO. The highest CO yield (3.51%) was found at the lowest H2O2 concentration. No methane was detected. For each of the organic compounds, the addition of H2O2 lowered the consumption of nitrate. However, the nitrate removal was higher than would be expected if the rate of organic oxidation by hydrogen peroxide was assumed to be much faster than the oxidation by nitrate. The discrepancy between the calculated (based on this hypothesis) and experimental results was biggest for phenol, indicating that the phenol oxidation by nitrate occurs fast. Nitrate and hydrogen peroxide are competitors for the oxidation of the organic compounds. Discussion: Proposed Reaction Mechanisms
Figure 11. Gaseous nitrogen products in the reaction of phenol with ammonium nitrate.
N2O and NO production decreased with increasing carbon concentrations. The NO production was the greatest for the oxidation of acetic acid and minimal for the phenol oxidation. Metals analysis for these experiments yielded concentrations no greater than 0.08 ppm for Cr, 0.24 ppm for Fe, 0.02 ppm for Mo, 0.07 ppm for Ni, 0.08 ppm for Pd, and 0.096 ppm for Ti. Reactions of Methanol, Acetic Acid, and Phenol with Ammonium Nitrate/Hydrogen Peroxide Mixtures. As discussed above, the oxidation of the organic compounds by ammonium nitrate resulted in the formation of NO. In an effort to eliminate NO emission, the oxidation of the organic compounds was investigated using a co-oxidant system of hydrogen peroxide with nitrate. The reactions of ammonium nitrate/hydrogen peroxide mixtures with methanol, acetic acid, and phenol were investigated at temperatures and pressures near 500 °C and 345 bar, respectively. Reactor residence times were approximately 35 s. The experimental conditions and results for the experiments are listed in
Ammonium Nitrate System. The main reaction pathway for NH4NO3/organic compounds was proposed (Figure 13) and will be discussed in detail in this section. Thermal decomposition of neat ammonium nitrate has been the subject of investigation for many years (Federoff, 1960; Medard, 1989). Reaction products, usually N2O, N2, and NOx, depend significantly on the reaction conditions (temperature, pressure, and contact of the samples). A free-radical mechanism was proposed for high-temperature reactions (>290 °C) (Rosser et al., 1963), while an ionic mechanism was proposed at relative low temperatures (Brower et al., 1989; Mearns and Ofosu-Asiedu, 1984). Related work in supercritical water includes the oxidation of ammonia by oxygen or hydrogen peroxide. In these systems, very low conversions were achieved even at high temperatures (>550 °C) (Killilea et al., 1992). The oxidation of ammonia by nitrate, however, was found to be fast, and high conversions were achieved under hydrothermal conditions (Luan et al., 1995; Bowman and Fulton, 1995). The reaction chemistry of alkali nitrates was studied intensively by Dell’Orco (1994). NO2 was suggested to be the primary reactive species, formed through homolysis of HNO3. A detailed mechanistic investigation of the reaction of ammonia with nitrate is not currently available.
1564 Ind. Eng. Chem. Res., Vol. 36, No. 5, 1997 Table 3
CH3OH
CH3COOH
C6H5OH
temp (°C)
pressure (bar)
RT (s)
[C]0 (M)
[NO3]0 (M)
[NH3]0 (M)
[H2O2]0 (M)
N 2O yield (%)
N2 yield (%)
N balance (%)
C balance (%)
NO3 removal (%)
487 491 491 492 496 500 500 500 496 496 496 498
349.0 340.6 343.9 342.3 339.5 345.7 339.2 344.4 344.8 344.8 344.8 342.4
35 33 33 33 35 37 35 37 36 35 33 31
0.41 0.37 0.39 0.35 0.40 0.39 0.38 0.37 0.36 0.30 0.34 0.33
1.10 1.13 1.10 0.92 1.00 0.95 0.96 0.98 1.12 1.09 1.09 0.97
0.98 0.98 0.93 0.83 1.00 0.97 0.96 0.99 0.99 1.01 0.95 0.94
1.20 0.90 0.60 0.30 0.90 0.60 0.30 0.15 1.20 0.90 0.60 0.30
19.01 17.47 11.77 10.66 25.41 20.45 17.46 16.61 17.26 16.90 17.34 13.48
59.15 58.54 69.59 98.26 68.19 65.40 63.44 71.44 55.15 53.39 62.83 66.11
96.08 91.25 89.01 109.03 105.43 96.10 84.17 88.64 89.36 87.08 94.73 87.05
74.83 90.85 86.39 90.64 105.43 106.74 92.28 92.68 82.00 77.05 85.85 86.74
66.29 71.60 87.97 99.89 76.66 82.30 95.80 100.00 68.15 67.87 73.05 86.16
the intermediate •NH2 radical is expected (reaction 4). Subsequently, reactions of •NH2 with NO (reaction 5) or NO2 (reaction 6) produce ONNH2 and O2NNH2, which quickly decompose to N2 or N2O, respectively. The ratio of N2/N2O is determined by the NO/NO2 ratio based on the proposed mechanism. This proposed reaction mechanism would result in the global equations for the formation of N2 and N2O as follows:
NH3 + HNO3 f N2O + 2H2O 5NH3 + 3HNO3 f 4N2 + 9H2O Figure 12. NO yields for the oxidation of the organic compounds by ammonium nitrate and hydrogen peroxide.
Given that N2/N2O is 4:1, the global reaction would be
6NH3 + 4HNO3 f 4N2 + N2O + 11H2O
Figure 13. Proposed reaction mechanism.
Reactions in this study were carried out near 500 °C and 345 bar with a water density of 0.14 g/cm3. Both N2 and N2O were observed as reaction products at a ratio of about 4:1. Under these conditions, water behaves as a nonpolar solvent; therefore, free-radical reactions could be dominant. On the basis of the experimental results and freeradical mechanism by Brower and co-workers and the work of Dell’Orco, we propose the formation of N2 and N2O, in the absence of organic compounds, through the following pathways (Figure 13): At higher temperatures and lower pressures, ammonium nitrate dissociates into ammonia and nitric acid in supercritical water (reaction 1). Homolysis of nitric acid produces reactive hydroxyl radical •OH and NO2 (reaction 2). This initial step would be the ratedetermining step. The hydroxyl radical is highly reactive. Hydrogen abstraction of ammonia by •OH to form
The stoichiometric ratio is predicted to be HNO3/NH3 ) 0.67. This is consistent with the experimental result of 0.65 determined by Figure 2. The equilibrium in the system NO-NO4-N2O4-N2O3-O2 has been previously studied in the gas phase and shown to be temperature dependent (Smith and Missen, 1982). At temperatures below 177 °C, NO2 is the major component (>99%). As temperature increases, more and more NO is formed. Besides the NO/NO2 equilibrium, minor side reactions could affect the production of N2 and N2O. For example, NO2 could react with NH3, producing NO which would enhance the N2 yield. Contributions from this reaction should be very small since •OH radicals are much more reactive than NO2. This assumption is supported by the fact that in the gas phase the reaction of NO2/NH3 is reported to be relatively slow (Bedford and Thomas, 1972). A second side reaction may come from the effect of water on the reaction mechanism. Reactions of reactive intermediates with H2O should be considered. A third possible side reaction is the decomposition of N2O to N2 which has been observed in the gas phase reactions at high temperatures. However, N2O decomposition is not likely at temperatures near 500 °C. This is supported by the fact that the observed N2/N2O ratio did not increase with increased residence time (Table 1). This reaction mechanism does not predict the production of NOx. If the formation of •NH2 and NOx is balanced, sufficient •NH2 is formed to consume all the NOx. Ammonium Nitrate/Organic Compound System. Addition of organics introduces the competition for •OH radical. According to eq 2, equal amounts of •OH and NOx are produced in the homolysis of nitric acid. When organics are present, some of the •OH will react with the organics in the reaction mixture. Oxidation of organic compounds has been shown to undergo
Ind. Eng. Chem. Res., Vol. 36, No. 5, 1997 1565
radical reactions with •OH (Atkinson, 1986). Hydrogen abstraction by reactive radicals •OH was suggested as the initial step (reaction 7). Equation 7 is expected to be fast and in competition with eq 4. There will not be enough •NH2 to react with all the NOx produced. Indeed, NO was observed from the reaction of all three organic compounds in the gas phase by IR spectroscopy. Reactions of organics with NO were reported to be slow in the gas phase (Bedford and Thomas, 1972; Odenbrand et al., 1986). This is consistent with our experimental results. We have observed NO production in the presence of excess organics. The NO production decreased with increased initial carbon concentration. Little NO gas was detected in the presence of a large excess of organics (Figures 6-8). This suggests that NO/organic reactions do occur at a significant rate under our experimental conditions. As was discussed earlier in the section on ammonium nitrate reactions, the N2/N2O ratio is determined by the NO/NO2 ratio (about 4). This is only true when no other reaction pathway exists. With organics/NH4NO3 reactions, we observed a N2/N2O ratio of about 4 when excess nitrate was available. In the presence of excess organics, the N2/N2O ratio increased with increasing reducing equivalent (i.e., organic carbon concentration) in the feed solution (Figures 9-11). This might indicate that the organics react more rapidly with NO2 than with NO, thus increasing the NO/NO2 ratio. CO and H2 were observed in the organics/NH4NO3 experiments run with a substoichiometric concentration of oxidizer (Figures 6-8). The CO production increased with increasing initial carbon concentration. CH2O has been proposed as a key intermediate in the formation of CO (Helling and Tester, 1987). CH2O is also mentioned by Boock and Klein (1993) in their proposed reaction mechanism for methanol and acetic acid oxidation by O2.
H2CO f H2 + CO
(8)
Equation 8 suggests that equal amounts of CO and H2 should be produced. This is true for acetic acid, while for methanol, hydrogen production is slightly higher than CO. The water-gas shift reaction, which was previously observed in supercritical water (Helling and Tester, 1988), will also effect the CO and H2 concentrations.
CO + H2O f H2 + CO2
(9)
This reaction may be affected by the solution pH. Under basic conditions, CO2 exists as HCO3- or CO32-, pushing the reaction equilibrium to the right and decreasing the CO concentration but increasing the H2 concentration. Addition of Hydrogen Peroxide. In aqueous solutions, hydrogen peroxide decomposes into water and oxygen as follows:
2H2O2 f 2H2O + O2 Lin and co-workers (1991) concluded that the initial step of H2O2 decomposition in water is not a simple homolytic cleavage if no ionic catalysts are present in the solution. However, homogeneous catalytic effects, due to the presence of metallic ions (Cu+ in ppb concentrations) were suggested to enhance the formation of •OH radicals in the decomposition of H2O2. Since concentrations of metal ions were present in the experi-
ments of this work, hydrogen peroxide probably functioned as a radical initiator as well as an oxygen source. As was mentioned before, NO production decreased with increasing amounts of H2O2 in the feed solution, but stoichiometric amounts of H2O2 were needed to have no detectable NO concentrations in the gaseous effluent. When hydrogen peroxide provides all the oxygen needed for the organic destruction, reaction 7 will become less dominant and equal amounts of •NH2 and NOx will be produced, resulting in complete reaction of all NOx. Conclusion Ammonium nitrate is an effective oxidizer for the tested organic compounds. Reactions were completed within less than 30 s at 500 °C and 5000 psi. Oxidation was only limited by the availability of the oxidizing agent. Major reaction products were N2, N2O, NO, and CO2 (bicarbonate and carbonate). Small amounts of CO, H2, and residual NH3 and TOC were detected in the effluent for higher than stoichiometric amounts of organics. Residual nitrate and small amounts of nitrite existed in the effluent when less than stoichiometric organic was present. NO was formed at less than 8% (of the total N in the feed solution) and decreased with increasing TOC in the feed. The relative oxidation rate of ammonia by nitrate was comparable to that of phenol and much faster than that of methanol and acetic acid. The phenol oxidation showed distinctly different chemistry from methanol and acetic acid. The addition of hydrogen peroxide as a co-oxidant still produced NOx at substoichiometric H2O2 to carbon concentrations. It, therefore, does not offer a cheap solution for the NOx problem. To eliminate the presence of NOx in the gas effluent, ammonium nitrate could be added in substoichiometric concentrations to the carbon in the waste stream. As was discussed before, at excess carbon concentrations, no NOx will be detected, but the effluent will still contain high carbon levels. A polishing step with oxygen or hydrogen peroxide will easily destroy the remaining carbon waste. Acknowledgment This research was funded by the U.S. Army Corps of Engineers Research Laboratory. The help of Rhonda McInroy with analytical analysis is greatly appreciated. Literature Cited Atkinson, R. Kinetics and Mechanisms of the Gas-phase Reactions of the Hydroxyl Radical with Organic Compounds. Chem. Rev. 1986, 86, 69. Bedford, G.; Thomas, J. H. Reaction between Ammonia and Nitrogen Dioxide. J. Chem. Soc., Faraday Trans. 1 1972, 2163. Boock, L. T.; Klein, M. T. Lumping Strategy for Modeling the Oxidation of C1-C3 Alcohols and Acetic Acid in High Temperature Water. Ind. Eng. Chem. Res. 1993, 32, 2464. Bowman, L. E.; Fulton, J. L. Hydrothermal Oxidation of Ammonium by Nitrate: a Raman Spectroscopy Study. In Physical Chemistry of Aqueous Systems: Proceedings of the 12th International Conference on the Properties of Water and Steam; White, H. J., Sengers, J. V., Neumann, D. B., Bellows, J. C., Eds.; Begell House: New York, 1995. Brower, K. R.; Oxley, C. O.; Mohan, T. Evidence for Homolytic Decomposition of Ammonium Nitrate at High Temperature. J. Phys. Chem. 1989, 93, 4029. Cox, J. L.; Hallen, R. T.; Lilga, M. A. Thermochemical Nitrate Reduction. Environ. Sci. Technol. 1992, 28, 423.
1566 Ind. Eng. Chem. Res., Vol. 36, No. 5, 1997 Dell’Orco, P. C. Reactions of Inorganic Nitrogen Species in Supercritical Water. Ph.D. Dissertation, The University of Texas at Austin, Austin, TX, 1994. Federoff, B. T. Encyclopedia of Explosives and Related Items; Picatinny Arsenal: Dover, NJ, 1960; Vol. I, p A311. Foy, B. R.; Dell’Orco, P. C.; Breshears, D.; Buelow, S. J.; Ely, J.; Funk, K.; Le, L. A.; McInroy, R. E.; Oldenborg, R. C.; Robinson, J. M.; Sedillo, M.; Wilmanns, E. Hydrothermal Kinetics of Organic and Nitrate/Nitrite Destruction for Hanford Waste Simulant. Los Alamos Unclassified Report LA-UR-94:3174, 1994. Haar, L.; Gallaghe, J. S.; Kell, G. S. NBS/NRC Steam Tables; Hemisphere Publishing Corp.: Washington, DC, 1984. Harradine, D. M.; Buelow, S.; Dellorco, P. C.; Dyer, R. B.; Foy, B. R.; Robinson, J. M.; Sanchez, J. T.; Spontarelli, Wander, J. D. Oxidation Chemistry of Energetic Materials in Supercritical Water. Hazard. Waste Hazard. Mater. 1993, 10, 233. Helling, R. K.; Tester, J. W. Oxidation Kinetics of Carbon Monoxide in Supercritical Water. Energy Fuels 1987, 1, 417. Helling, R. K.; Tester, J. W. Oxidation of Simple compounds and Mixtures in Supercritical Water: Carbon Monoxide, Ammonia, and Ethanol. Environ. Sci. Technol. 1988, 22, 1319. Killilea, W. R.; Swallow, K. C.; Hong, G. T. The Fate of Nitrogen in Supercritical Water Oxidation. J. Supercrit. Fluids 1992, 5, 72. Lee, D. S.; Gloyna, E. F. Efficiency of H2O2 and O2 in Supercritical Water Oxidation of 2,4-Dichlorophenol and Acetic Acid. J. Supercrit. Fluids 1990, 3, 249. Lin, C. C.; Smith, F. R.; Ichikawa, N.; Baba, T.; Itow, M. Decomposition of Hydrogen Peroxide in Aqueous Solutions at Elevated Temperatures. Int. J. Chem. Kinet. 1991, 23, 971. Luan, L.; Proesmans, P. I.; Foy, B. R.; Buelow, S. J. Hydrothermal Oxidation of Ammonia/Organic Waste Mixtures. Los Alamos Unclassified Report LA-UR-95-815, 1995.
Mearns, A. M.; Ofosu-Asiedu, K. Ammonium Nitrate Formation in Low Concentration Mixtures of Oxides of Nitrogen and Ammonia. J. Chem. Technol. Biotechnol. 1984, 43A, 350. Medard, L. A. Accidental Explosions; Wiley: New York, 1989; Vol. II, p 545. Odenbrand, C. U.; Andersson, L. A.; Brandin, J. G.; Lundin, S. T. Catalytic Reduction of Nitrogen Oxides. Appl. Catal. 1986, 27, 363. Rosser, W. A.; Inami, S. H.; Wise, H. The Kinetics of Decomposition of Liquid Ammonium Nitrate. J. Phys. Chem. 1963, 67, 1753. Smith, W. R.; Missen, R. W. Chemical Reaction Equilibrium Analysis; Wiley: New York, 1982. Tester, J. W.; Holgate, H. R.; Armellini, F. J.; Webley, P. A.; Killilea, W. R.; Hong, G. T.; Barner, H. E. Supercritical Water Oxidation Technology: A Review of Process Development and Fundamental Research. In Emerging Technologies in Hazardous Waste Management III; Tedder, D. W., Pohland, F. G., Eds.; ACS Symposium Series 518; American Chemical Society: Washington, DC, 1992; Chapter 3. Webley, P. A.; Tester, J. W.; Holgate, H. R. Oxidation Kinetics of Ammonia and Ammonia-Methanol Mixtures in Supercritical Water in the Temperature Range 5300-700 °C at 246 bar. Ind. Eng. Chem. Res. 1991, 30, 1745.
Received for review March 27, 1996 Revised manuscript received December 2, 1996 Accepted December 6, 1996X IE9601716
X Abstract published in Advance ACS Abstracts, January 15, 1997.
