uniform, which must affect the quality of the product; as a consequence, mild reaction parameters must be chosen to ensure the quality of the product. I n this latter case, the capacity of the plant is small, as proved by the productivity of the reactor. Only a schematic outline of the Atlas Powder process is known; however, on the basis of this outline, the authors assume that the parameters of the process are similar to those described in the present work. T h e difference in the two processes is in the gas circulation system-Le., in the Atlas Powder process, the gas is recirculated, while the Hungarian process uses a single-pass gas flow. T h e advantage of the authors' process is that with the aid of stepwise heating and gradual conversion, a product of high and uniform quality may be produced with a high capacity. The lack of hydrogen recirculation is a novel solution. This was possible because the plant was located near a n ammonia plant, which is a high capacity hydrogen source. T h a t a synthesis gas of 25% Nz by volume was used for hydrogenation rather than pure hydrogen is novel. I n using the 420 ton-per-day ammonia synthesis plant as a hydrogen source for 1000 tons of sorbitol per year, the flow of the gas is modified as follows (volumes are expressed under standard conditions). T h e fresh synthesis gas (0.32%) is diverted through the sorbitol plant, reduced in volume by the amount of hydrogen necessary for the reduction, and expanded to a lower pressure (20 atm.). Following the absorption of carbon dioxide, it is led back to the main flow of the synthesis gas. T h e 7570 Hz content of the synthesis gas is reduced by the hydrogen consumption of the sorbitol plant by only 0.05%. The above data show that the operation of the ammonia synthesis plant is not disturbed. Factors which influence the operation more significantly include compression capacity change caused by the winter-summer temperature differences, fluctuations of the efficiency of the gas purification, and fluctuation of the efficiency of hydrocarbon reforming. If the sorbitol production were increased tenfold-Le., to 10,000 tons per year-the gas diverted to the sorbitol plant could be compensated for by production in the ammonia synthesis plant of a gas of greater than 75% hydrogen content.
Hydrogenation by the aid of ammonia synthesis gas provides significant economical advantages. The investment and production costsfor a plant producing 1000 tons of sorbitol per year by electrolytically produced hydrogen and an independent gas circulation, and sorbitol production by hydrogen diverted from the 420 ton-per-day ammonia synthesis plant are compared below. Investment Costs
A.
B.
Investment costs for 1 std. cu. meter per year synthesis gas compressed to 300 atm. Investment costs for 1 std. cu. meter pure HZ in synthesis gas a t 300 atm. Investment costs for 1 std. cu. meter per year electrolytic hydrogen compresses to 300 atm.
$0.013 0.017 0.15
Production Costs
A.
B.
Hydrogenation with synthesis gas Production costs of synthesis gas necessary to produce 1 ton of sorbitol Hydrogenation with electrolytic hydrogen Production costs of hydrogen necessary to produce 1 ton of sorbitol
$ 4.40
$20.80
Thus, combination of the sorbitol with the ammonia plant lowers the investment cost of the production of 1 ton of sorbitol by $24 and the production costs by $16.40, as compared with the same costs using an independent hydrogen source. For a plant with a capacity of 1000 tons per year, the savings are $16,400 and $24,000 per year for production and investment costs, respectively. literature Cited
Ashcroft, W. K., Chem. Age (London) 89, 907 (1963). Brahme, P. H., Pai, M. V., Narsimhan, G., Brit. Chem. Eng. 9, 684 (1964). Chem. Eng. 59, 208 (1959). Fedor, W. S., Ind. Eng.Chem. 52, 282 (1960). Haidegger, E., Magyar Ke'm. Lapis 10, 351 (1961). Hoffman, H., Bill, W. I., Chem. Zng. Tech. 31, 81 (1959). Robertson, J. H., Ind. Chemist 39, 233 (1966). RECEIVED for review March 8, 1966 ACCEPTEDAugust 29, 1967
HYDROGEN PEROXIDE CATALYTIC OXIDATION OF REFRACTORY ORGANICS IN MUNICIPAL WASTE WATERS D. F . B I S H O P , G . S T E R N , M . F L E I S C H M A N , A N D L . S . M A R S H A L L Advanced Waste Treatment Research Activities, Cincinnati Water Research Laboratory, Federal Water Pollution Control Administration, Cincinnati, Ohio 45226
organic pollutants entering the nation's water reoxidation in municipal waste treatment facilities or in natural water courses. These refractory organics, after chlorination, may form noxious chlorinated organics and cause general deterioration in water quality. T h e increasing re-use of water demands better removal of organic contaminants. ANY
M sources are not effectively removed by biological
110
I&EC PROCESS DESIGN A N D DEVELOPMENT
Therefore, an experimental study was undertaken to determine the technical feasibility of combining active oxygen (hydroxyl radicals) chemical oxidation and molecular oxygen oxidation (autoxidation) to destroy the broad spectrum of refractory organics in municipal waste waters. An important aim of the study was to determine in various waste waters the capabilities of autoxidation initiated with small amounts of
Hydrogen peroxide catalytic oxidation of organic residuals in municipal waste waters involves a free radical ( - O H ) oxidation and is effective only in a 3 to 5 pH range. Both ferric and ferrous salts are efficient catalysts, but the ferric system requires elevated temperature. Oxidation efficiencies of approximately 60% are achieved for stoichiometric charges of hydrogen peroxide to chemical oxygen demand (COD) of the waste water. The oxidation process preferentially attacks alkyl benzene sulfonate with 98% removal, and oxidizes from 30 to 65% of the total organic carbon to carbon dioxide. It also reduces the C O D of water from carbon treatment systems by 50 to 65% of the initial COD. Autoxidation occurs in the oxidation process, but with average chain lengths of less than six oxidation events per initiation event. A more economical source of free radical ( * O H ) without a restricted p H range is needed for practical application.
the hydroxyl radical. T h e ferric ion--hydrogen peroxide and ferrous ion-hydrogen peroxide catalytic systems were selected as convenient laboratory sources of the hydrovyl radical. Hydrogen peroxide and iron salts separately are not effective oxidants of the complex organic materials in municipal waste waters. When the hydrogen peroxide and iron salts are combined, however, the peroxide decomposes (Barb et al., 1959) to produce the hydroxyl radical oxidant.
+ Fe+3 + H 2 0 2
kl
Fe+2
+ OH- f . O H Fe+2 + HOZ. + H +
H202 + Fe+2 k2
+
(1 1
(2)
When ferrous salts are used, the hydroxyl radical is produced immediately by the rapid reaction between ferrous ion and hydrogen peroxide (Equation 1, rate constant k l = 4.45 X 10s e - 9 4 0 0 / R 1 moles liter-' second-'). With ferric salts, the hydroxyl radical is produced by a two-stage process with the slow reaction between ferric ion and hydrogen peroxide (Equation 2, k 2 = 1.1 X 1 0 * 4 e - z * * 0 0 0 / R 1 moles liter-' second-') followed by the rapid reaction between the produced ferrous ion and additional hydrogen peroxide. Other reactions (Barb et al., 1959) in the decomposition mechanism compete with the organic substrate for the available active oxidant in the hydrogen peroxide. Fe+2
+ -OH
-+
HOn. % 0
+ 02Hz02 + .OH Fe+3
Fef3
2 -
f H+
+ O2 HOz. + HzO
-+
+
+ OH-
Fe+2
(3)
(4)
(5) (6)
Thus a n unknown part of the theoretical oxidation equivalent in hydrogen peroxide is actually available for oxidation. The final decomposition products of the hydrogen peroxide catalytic systems are water, oxygen, and ferric hydroxide. Thus, these systems do not contribute soluble pollutants to the waste water. I n addition, the ferric hydroxide acts as a flocculant which removes suspended solids from the waste water. I n the complex organic-inorganic waste water substrate, many reactions between the active oxidant and the substrate can occur. T h e hydroxyl radical has a large electron affinity (136 kcal.) and not only oxidizes all organic compounds containing hydrogen, but also reacts with most halide (X-) ions (Cl-, Br-) to produce X . radicals (Uri, 1952). I n a simplified general oxidation mechanism for hydroxyl radicals and molecular oxygen oxidation suggested by Uri (1961), a hydroxyl radical abstracts a hydrogen atom from the organic substrate, producing a n organic radical (R .).
RH
+ .OH
+
R.
+ H20
(7)
T h e organic radical ( R . ) reacts rapidly with environmental molecular oxygen to form a n organic peroxy radical ( R O O . ) .
R.
+
2 +
0
ROO.
(8)
T h e organic peroxy radical abstracts a hydrogen atom from the organic substrate, producing an organic hydroperoxide (ROOH) and another organic radical.
ROO.
+ RH
+
ROOH
+ Re
(9)
I n the presence of iron salts, the decomposition of hydroperoxide produces chain-continuing radicals and accelerates the over-all autoxidation rate.
+ Fe+2 ROOH + Fe+3
ROOH
-+
+ Fe+3 + OHROz. f Fe+2 + H +
RO.
+
(10)
(11)
I n the complex substrate in waste waters, the simplified reaction mechanism is undoubtedly not the only type of oxidation reaction in the system. Nonchain hydroxyl radical oxidation, direct hydrogen peroxide oxidation, and oxidation by X . radicals, produced from X- reactions with hydroxyl radicals, probably also occur. Experimental
Scope. T h e oxidation studies were conducted on refractory municipal wastes in secondary effluents, concentrated municipal effluents from distillation processes, collapsed foamate from foaming studies, and very dilute municipal effluents from carbon absorption treatment. T h e chemical oxygen demand (COD) of the test effluents varied from approximately 10 to 2100 mg. per liter. Most of the tests were performed on filtered (Whatman No. 5 filter paper) municipal secondary effluents with COD'S varying from 30 to 150 mg. per liter. T h e secondary effluents included waste waters of chiefly domestic origin and waste waters of mixed industrial and domestic content. T h e experimental variables were concentration and nature of the refractory organics, concentration of hydrogen peroxide, concentration of iron salt (catalyst), temperature, p H of the system, oxygen pressure, and reaction time. Three catalytic oxidation systems, distinguished by rate factors, were studied. They are ferric ion-hydrogen peroxide, ferrous ion-hydrogen peroxide, and ferrous ion-ferric ionhydrogen peroxide systems. T h e ferric ion-hydrogen peroxide and ferrous ion-ferric ion-hydrogen peroxide systems required elevated temperatures for practical rates of reaction. Ferrous ion-hydrogen peroxide oxidation rates were practical a t ambient temperature. Since oxidation rates depend upon hydrogen peroxide decomposition rates, the hydrogen peroxide decomposition rate a t 65.8' C. with ferric salts was studied in distilled water and waste water over a 1.8 to 9 p H range. Rate studies of the rapid ferrous-peroxide reaction were not included. Analytical Methods. T h e parameters used to evaluate the oxidation processes were (mg. per liter) COD, hydrogen peroxide concentration, organic carbon concentration, and VOL. 7
NO. 1
JANUARY 1968
111
apparent alkyl benzene sulfonate (ABS) concentration. T h e C O D test, performed with duplicate samples, was the modified procedure (Dobbs and Williams, 1963) employing mercuric sulfate to eliminate chloride interference. In this study, the standard deviation of 192 duplicate COD'S on secondary waste water effluents was 1.9 mg. of oxygen per liter. The initial hydrogen peroxide charge was determined by direct pGtassium permanganate titration of the stock hydrogen peroxide solutions immediately before use. The hydrogen peroxide concentrations in the oxidation system were measured by the titanic sulfate spectrophotometric procedure (Eisenberg, 1943). Ferric ions in the concentrations used did not interfere in the spectrophotometric peroxide analysis. The organic carbon concentration was measured with a carbon analyzer (Beckman Carbonaceous Analyzer, COz analyzer calibrated in milligrams per liter of total carbon) (Van Hall et al., 1963). Acidification and nitrogen stripping removed the inorganic carbonates before analysis. The apparent ABS concentration was determined by the standard methylene blue procedure (American Public Health Association, 1960). The changes in COD and hydrogen peroxide concentration were combined to determine the oxidation efficiency (O.E.) of the oxidation process.
O.E. =
A COD (mg./literj
available 0
2
(mg./liter)
x
100
(12)
T h e available oxygen was the theoretical amount of reactive oxygen in the charged hydrogen peroxide based on two oxygen equivalents per mole of peroxide. Thus the oxidation efficiency was the ratio of the amount of oxidation of the waste water substrate as measured by its change in C O D to the maximum amount of oxidation possible from the hydrogen peroxide. Changes in the organic carbon concentrations of waste waters indicated the amount of organic carbon oxidized to carbon dioxide. In most tests, the charged hydrogen peroxide completely decomposed; however, in a few tests residual hydrogen peroxide, which interfered in the C O D analysis, was removed by adding appropriate amounts of sodium sulfite or titanous sulfate reducing agents. The amount of sodium sulfite needed to remove all the hydrogen peroxide was based on the peroxide concentration ascertained by means of an external titanic sulfate spot plate test. If excess of sulfite was added to the sample, the resulting ferrous ions produced by the reaction between excess sulfite and ferric ions were measured by the 2,2'-bipyridine spectrophotometric determination (Moss and Mellon, 1942) of ferrous ion. Ammonium fluoride was added to eliminate the interference of ferric ion in the ferrous determination. The appropriate correction for the presence of ferrous ions \+as then made on the experimental COD. Residual peroxide was also removed by rapidly titrating to the stoichiometric end point with a 2 7 ( titanous sulfate solution using the titanic sulfate-hydrogen peroxide color complex as its own internal indicator. The effect of flocculation by ferric hydroxide on the C O D reduction was eliminated by dissolving the ferric hydroxide with sulfuric acid or by using samples drawn from well mixed systems before performing the C O D analysis. Materials and Apparatus. ACS reagent grade ferric sulfate, ferric ammonium sulfate, and ferrous ammonium sulfate were used as catalysts. Preliminary laboratory tests revealed that below p H 5, the ammonium ion in the waste water or from the catalyst was not oxidized by the oxidation process. Thirty per cent ACS reagent grade hydrogen peroxide was used as the oxidant source. Research grade 0 2 and NZ were used to control the type of atmosphere. The specific resistance of the laboratory-demineralized water varied between 5 X l o 5 and 9 X 105 ohms. Acrylamide monomer (AMD), used for polymerization detection of free radicals, was supplied by the American Cyanamid Co. The initial screening tests were conducted in borosilicate glass beakers covered with aluminum foil. The waste water sample and the chemical reagents were mixed in the beakers and left a t ambient temperature or placed in a laboratory oven set a t 65' C . The apparatus for kinetic studies consisted of a 2-liter, threenecked distillation flask with 24/40 ground-glass joints, a 112
l&EC PROCESS DESIGN A N D DEVELOPMENT
variable-speed mixer with Teflon paddles, a Beckman expanded-scale p H meter with combination probe, and a NBS thermometer calibrated in 0.2" C. The apparatus was placed in a constant temperature bath that could be maintained to zk0.1" C. The stirrer was set at a constant 200 r.p.m. If atmospheric control was desired, oxygen or nitrogen was bubbled into the flask through a medium-pore glass sparger. A rotatable pressure vessel was constructed for use in the O 2 pressure studied a t temperatures up to 65" C. T h e 3-liter stainless steel vessel included provisions for pressuring the system and removing gas and liquid samples without losing the pressure in the vessel. T h e vessel was placed in a laboratory oven for elevated-temperature tests. Procedure. In the peroxide decomposition study, the ferric salt was added to the distilled water or waste water a t 65' C. in the 2-liter kinetic vessel. T h e pH was adjusted to the selected value, and the hydrogen peroxide was then added to the system. Samples of well-mixed ferric ion-hydrogen peroxide solution were periodically pipetted from the flask and mixed with titanic sulfate reagent. The decreasing concentration of peroxide was determined by measuring the light absorption of the titanic-hydrogen peroxide colored complex a t 420 mp. In the ferric ion-hydrogen peroxide oxidation system, the ferric salt was added to the waste water a t 65" C . and the p H adjusted to the appropriate value. A known amount of peroxide was added and the oxidation allowed to proceed for the selected reaction time. Appropriate analyses on the waste water were performed before, during, and after oxidation treatment. I n the ferrous ion-ferric ion-hydrogen peroxide system, the initial charge of hydrogen peroxide was added to the waste water a t 65" C., and the system's p H was adjusted to the desired value. Ferrous ion in less than stoichiometric amounts (Equation 1) was then added to the system. After the rapid ferrous ion-peroxide reaction (Equation 1) and the hydrolysis of the resulting ferric ions were completed, the p H was readjusted to the appropriate value, and the residual peroxide was decomposed by the ferric mechanism. In the ferrous ion-hydrogen peroxide tests. both a t ambient temperature and 65" C., the hydrogen peroxide was added to the waste water at the selected pH. Small aliquots of ferrous solution were added periodically to the hydrogen peroxidewaste water system over a 15- to 20-minute interval until all of the peroxide decomposed. Readjusting the p H of the system after each ferrous aliquot prevented large p H changes. In the polymerization study to verify the presence and availability of the free radicals, 10% by weight of acrylamide monomer was added to the waste water or distilled water under nitrogen atmosphere at either 65" C. or ambient temperatures. The hydrogen peroxide and ferrous or ferric salts were added to produce the free radicals and thus the polymerization. In the tests, the nitrogen gas was continuously bubbled through the 2-liter, three-necked vessels to prevent oxygen quenching of the polymerization chain. Specific autoxidation tests were conducted a t 65" C . under 1 atm. or 75 to 80 p.s.i.g. of molecular oxygen. The ferric ion-hydrogen peroxide oxidant was added to waste water either in a 2-liter flask through which oxygen was continuously bubbled or in the pressure vessel which was then pressurized with oxygen to 75 to 80 p.s.i.g.
Ferric Ion-Hydrogen Peroxide Decomposition. T h e ferric-peroxide decomposition system was studied over a 1.8 to 9.2 p H range a t 65.8" C. with a fixed total ferric charge of 0.001264 equivalent of ferric ion per liter of distilled water. Above p H 2.5, the system was heterogeneous, with ferric hydroxide forming the solid phase. In the initial stages of the decomposition a t approximately 10 to 1 equivalents ratios of hydrogen peroxide to total ferric salt, the decomposition (Figure 1) a t any fixed p H was apparently first-order with hydrogen peroxide. I n waste waters, the initial hydrogen peroxide decomposition (Figure 2) appeared first-order, but over a shorter range of hydrogen peroxide- to-ferric equivalents ratios than in distilled water. The hydrogen peroxide decomposition rate in
10.0
I."
I
0
20
Figure 1 .
40
60 80 100 TIME IN MINUTES
120
140
160
HzOz decomposition at various p H
System. Heterogeneous above pH 2.5 Solid phase. Ferric hydroxide Total ferric concentrations. 0.001 2 6 4 eq. per liter Temperature. 65.8' C. Initial H?O?concentration. 0.01 1 2 5 to 0.01 2 6 7 eq. per liter
-
5%
5.0
0
N
U I \ 0
n
0 "
N
6 2.0
1.0
0
20
Figure 2.
40
60 80 100 TIME IN MINUTES
120
140
160
Hz0~ decomposition in waste water
System. Heterogeneous Solid phase. Ferric hydroxide Temperature. 65.8' C. At K = 6.29 X lO-%nin.-' Waste water. Loveland secondary Total ferric concentration. 0.002528 eq. per liter At K = 3 . 3 0 X l O - h i n . - ' Waste water. Hamilton secondary Total ferric concentrotion. 0.001 4 2 6 eq. per liter At K = 2 . 5 0 X 10-3min.-1 Waste water. Lovelond secondary Total ferric concentration. 0.001 2 7 2 eq. per liter
peroxide was then slowly decomposed by the ferric mechanism. T h e apparent first-order rate constant, k (Figure 3), for the ferric mechanism in the distilled water, and, therefore, the peroxide decomposition rate, exhibited a minimum in the 3 to 5 p H range. Stumm (1964) indicated that below a p H range of 3. the predominant soluble ferric species was hydrated ferric ion [ F e ( H ~ 0 ) 6 ] +and ~ in the 3 to 4 p H range, the hydrated ferric complex [ F e ( O H ) ( H 2 0 ) j ] + 2 . Above p H 4, the soluble species changed from [Fe(OH) (H20)j]+2 to [Fe(OH)a(H?0)4]+ and finally to the ferrite complex F e ( 0 H ) 4-. T h e total concentration of soluble ferric species continuously decreased as the p H increased until appreciable Fe(0H)d- began to form a t about p H 8. Thus from p H 1.8 to approximately 4, the peroxide decomposition rate decreased with decreasing soluble ferric species. Above pH 4, however, the rapid increase in the rate of hydrogen peroxide decomposition. occurring a t very lo\v (10-5 to 10-8 mole per liter) concentrations of soluble ferric species, indicated a change in the decomposition mechanism. probably to a basecatalyzed decomposition on the surface of the Fe(0H)a floc. Effect of pH. Early in the study, it became evident that the p H of the peroxide oxidation or polymerization systems was the most important variable in all systems. I n the oxidation process. the maximum organic removal as measured by the COD (Figure 4) and, thus, the maximum oxidation efficiency occurred in a 3 to 5 p H range. T h e COD removal decreased rapidly above p H 4 and below pH 3. Flocculation and settling of the ferric hydroxide accompanying the oxidation (Figure 4. Hamilton data) improved the organic removal from 1 5 to 307, over the removal by oxidation alone. .4lthough the p H dependency of the organic removal process \vas reduced by the ferric hydroxide flocculation, the maximum removal still occurred in the 3 to 5 p H range. Polymerization of 10%. by weight, of acrylamide monomer (AMD) in waste water and in distilled water not only verified the production of free radicals. thus supporting the proposed oxidation mechanism. but also showed p H dependency. T h e qualitative polymerization data (Table I) indicated that the best polymerization in waste water occurred a t p H 3 . 5 . Below p H 3, the polymerization was less effective and a t p H 7 and above, did not occur. T h e effective p H range of oxidation and polymerization
J
I
waste waters, in contrast to the decomposition rate in distilled water increased as the ratio of peroxide to total ferric ions decreased. I n the complex waste water substrate, the production of organic acids and other oxidation products and the possibilities for chelating and complexing ferric ions may have been responsible for the variation in the behavior of the decomposition rates. The specific factors, however, were not determined. I n decomposition systems with ferrous salts, the initial ferrous charge reacted rapidly with hydrogen peroxide to produce more than the stoichiometric (Equation 2) reduction in peroxide. The greater than stoichiometric reduction in peroxide was attributed to the reaction between the hydroxyl radical and the hydrogen peroxide (Equation 6). T h e residual
H2OpDECOMPOSITION CATALYST: Fez (SO,)
'$LO
2.0
Figure 3.
3.0
4:O 510 6.0 AVERAGE p H
7:O
8.0
S!O
Decomposition constant a t various p H
System. Heterogeneous above pH 2.5 Solid phase. Ferric hydroxide Total ferric concentrotion. 0.001 2 6 4 eq. per liter Temperature. 65.8' C. Initial HzOs concentration. 0.01 1 2 5 to 0.01 2 6 7 eq. per liter
VOL. 7
NO. 1
JANUARY
1968
113
-o-00
20
40
Below p H 3, more of the charged ferric ions remained in solution and were present as hydrated ferric ions. The larger number of soluble ferric ions in the solution below p H 3 accelerated not only the decomposition of the peroxide through Equations 1 and 2 but also the competitive reactions (Equations 3 to 6 ) , which reduced the number of hydroxyl radicals available for oxidation. Ferric species and ferric hydroxide are present in the ferrous-ferric-hydrogen peroxide systems, and their observed p H dependency occurred as in the ferricperoxide system. HAMILTON
60
- Fe+' '
'
80
Results I00
I20
AVERAGE O H
Figure 4.
Reduction of COD a t various pH
Solid phose. Ferric hydroxide Toto1 catalyst concentrotion. 0.001 264 eq. per liter Initio1 H202 concentrotion. 0.01 251 to 0.01 264 eq. per liter Temperature. 65' C. Reaction period. 22 hours
corresponded to the p H range in which the ferric-catalyzed hydrogen peroxide decomposition rate passed through its minimum. Although the effects of the various soluble ferric species are not fully understood, the slowest production of hydroxyl radicals occurred in the 3 to 4 p H range where is the predominant ferric species. [Fe(OH) (H20) The p H dependency of ferric-peroxide oxidation and polymerization processes was explained by the behavior of the ferric-catalyzed peroxide decomposition rate. Above p H 4, the rapid hydrogen peroxide decomposition, probably on the surface of the ferric hydroxide floc, would not produce appreciable amounts of available hydroxyl radicals in the solution. Since hydrogen peroxide decomposition ultimately produces molecular oxygen, the rapid decomposition above p H 4 not only reduced the available hydroxyl radicals in solution but also may have suppressed polymerization by the rapid production of molecular oxygen.
Oxidation. With all catalytic systems, residual organics as indicated by the C O D (Table 11) remained in municipal secondary waste water effluents after oxidation. As indicated by the reduction in the organic carbon concentration, between 36 and 66% of the organic load was converted to carbon dioxide. The oxidation efficiency increased as the ratio of charge oxidant to initial C O D in the waste water decreased. The residual organics exhibited increasing resistance to oxidation with increasing amounts of peroxide (Figure 5), probably related to decreasing organic concentrations as well as increasing chemical resistance to oxidation. Alkyl benzene sulfonate (ABS) was preferentially attacked by the oxidation process (Figure 6 ) . The apparent ABS concentration (methylene blue procedure) was reduced by 98% in the ferrous ion-ferric ion catalytic system even a t ambient temperature. For extremely small ferrous charges where negligible COD reductions occurred, 60% of the apparent ABS was destroyed. The ferric ion-peroxide system was tested on organic residuals remaining in waste water effluents after granular activated-carbon treatment. For stoichiometric oxidant charges, the oxidation (Table 111) of the organics, refractory to granular carbon treatment, was nearly as effective as the oxidation in secondary effluents. As in secondary effluents, the oxidation efficiency increased as the ratio of charge oxidant to initial COD decreased. The COD residual of 5 to 10 mg
Table I. Polymerization 107, by weight AMD solution purged by NZfor 0.5 hour Temp., ' C . System Initiator In it ial Max. PH Visual Polymer 65 80 Gel HzO Fe+*-HzOz 32 40 Gel HzO Fe +2-H20z Hamilton Fe +3-H202 65 75 Gel Hami 1ton Fe +2-Fe+3-Hz02 65 75 Gel Hamilton Fe+3-H20~ 66 73 2.40 Sirupy 66 69 3.57 Gel Hamilton Fe+3-Hz02 66 Hamilton Fe +3-H2 0 2 66 7.23 None 66 Hamilton Fe +3-H2O 2 66 None 11.75 a Oslwald oiscometer time before and after polymerization. * Too viscous to be measured in Ostwald viscometer. ~~
Time,a See. Before
After
30.6 30.6
b b
30.6
30.6 36.3
30.3
Table II. Oxidation and Organic Carbon Removal Available oxygen, 101.1 mg./liter; test pH, 3.8; temp., 65" C. Original Filtered COD, Final C O D , COD Secondary Mg./Liter Mg./Liter Reduction, H amiltona 79.8 36 .O 54.9 Lebanon 47.2 12.5 73.5 Loveland 126 62.0 50.8 55.4 20.0 Remington 63.9 a p H 3.5; available oxygen, 84.7 mg./iiter; catalyst, ferric ion.
114
l & E C PROCESS D E S I G N A N D DEVELOPMENT
70
Oxid. E$c., 70 51.7 34.3 63.3 35 . O
Organic Carbon, Mg./Liter Orig. Final 21.1 9.2 14.7 5 .O 35.1 22.5 15.8 7.5
%
Reduction 36.1 66.0 36.0 52.5
'
906 -
I
I
I
T
I
,
-
__
7 -
1
7 -
Iz
I
4 u >-
60:
w
