Environ. Sci. Technol. 1988, 22, 761-767
Destruction of Pollutants in Water with Ozone in Combination with Ultraviolet Radiation. 3. Photolysis of Aqueous Ozone Gary R. Peyton”,? SumX Corporation, 221 1 Denton Drive, Austin, Texas 78761
William H. Glaze$ Graduate Program in Environmental Sciences, University of Texas at Dallas, Richardson, Texas 75080
Photolytic ozonation (ozone/UV) has been known for some time to effectively destroy organic compounds which are refractory to ozonation. The mechanistic details, however, have not previously been elucidated. The present study demonstrates that the first step in the complex mechanism is the photolysis of aqueous ozone to produce hydrogen peroxide. This step is followed by secondary reactions which produce hydroxyl radical, the active species which is most responsible for organic compound destruction. In the presence of oxygen, reaction of many organic compounds with hydroxyl radical results in the production of superoxide and/or hydrogen peroxide, both of which may react further with ozone to produce more hydroxyl radicals.
Introduction This investigation was part of a broader study ( I ) , in which the ultimate objective was the pilot-scale evaluation of photolytic ozonation for the destruction of trihalomethane precursors in water. The first two papers (2,3) in this series described the destruction of tetrachloroethylene and of natural trihalomethane precursors by photolytic ozonation. The disappearance of those substances was described phenomenologically, and an empirical mathematical model was developed to express the rate of substrate destruction in terms of the ozone dose rate [in mg/(L.min)] and the ultraviolet light intensity (in W/L). It was recognized at that time that a detailed description of the kinetics and mechanisms of this complex process was needed. Therefore, an in-depth study was initiated to investigate those questions. Background As outlined in the first paper of this series (2),photolytic ozonation was developed in the early 1970s for the treatment of cyanide-containing wastes. It has been shown by several investigators (2,4-11) to be more effective for the destruction of some organic compounds than ozonation alone and has significant potential as a water treatment process for the destruction of organic compounds. Many authors have proposed (4-7) that hydroxyl or other free radicals are responsible for the ability of photolytic ozonation to destroy compounds which are refractory even to ozonation, but these hypotheses have not been supported directly by experimental data. Indeed, Leitus et al. (8) have given arguments against hydroxyl radical involvement and found photolytic ozonation to be only slightly more effective than ozonation alone in some cases. Some investigators (e.g., ref 12 and 13) have speculated that photolysis of aqueous ozone produces hydroxyl radical Present address: Aquatic Chemistry Section, Illinois State Water Survey, 2204 Griffith Drive, Champaign, IL 61820. Present address: Environmental Science and Engineering, School of Public Health, University of California, Los Angeles, CA 90024.
*
0013-936X/88/0922-0761~01.50/0
directly, analogous to the gas-phase reaction (14). Taube (IS),on the other hand, reported that ozone photolysis at 254 nm in aqueous acetic acid solution led to the production of a “stoichiometric amount” of hydrogen peroxide. As will be seen below, these two statements are not mutually exclusive. The present study was undertaken with the objective of determining whether ozone photolysis in aqueous solution results in the direct production of hydrogen peroxide (eq 1)or the direct production of hydroxyl radical (eq 2). These two candidate initiation steps will be referred to as case I and case 11, respectively.
O3+ H20 -% O2 + HzOz O3+ H 2 0 -% O2 + 2’OH
(2)
Experimental Section Ozone photolysis studies were carried out in a continuously sparged stirred tank photochemical reactor operated in the batch mode with respect to liquid. The apparatus used for the studies was similar to that described previously (1-3,10) and is shown in Figure 1. The 2-L stirred tank reactor was made from Pyrex pipe (6-in. i-d.), with top and bottom heads cut from 1/4 in. thick PTFE sheet. The reactor contained four baffles and a six-blade paddle-type impeller, made from aluminum, with the standard CSTR relative dimensions (16-18). All tubing and fittings were glass or PTFE. Four quartz UV lamp sheaths, open on both ends, extended vertically through the reactor. These were sealed into the reactor heads with viton O-rings in compression glands. All work reported in this paper was done with low-pressure mercury lamps, type G8T5, from American Ultraviolet (Chatham, NJ). These lamps are rated at 8-W power consumption with 1.6 W of UV output (at 100 h of life) at 254 nm. Ihtensities measured radiometrically, averaged cylindrically around the lamp, gave a UV power output 0.43 W per lamp at the time of the present experiments. Stirring speed was set at 550 rpm with a phototachometer. This speed was found by Prengle (16) and Yocum (17) to give good mass-transfer performTable I. Experimental Conditions for Photolytic Ozonation Runs parameter
value
ozone dose rate: mol/ (Lemin) 3.0 x UV dose rate,b einsteins/(L.min) 2.8 x 10-5 maximum mass transfer rate: mol/(L.min) 1.1 x IO“ maximum peroxide accumulation rate: mol/(L.min) 1.1x gas flow rate, SLM 0.2 reactor volume, L 2.0 stirring speed, rpm 550 aApplied. bLeaving lamp, i.e., not corrected for absorption by lamp wells, reactor walls, or reactor contents. ‘In the absence of ultraviolet irradiation; see Figure 5. dRun in 0.015 M acetic acid, pH 3.3; see Figure 3.
0 1988 American Chemical Society
Environ. Sci. Technol., Voi. 22, No. 7, 1988 781
~~~
~
Table 11. Oxidant Photolysis Experiments
CSTR
oxidant added ozone ozone ozone ozone HzOz
series I I1 I11 IV V
scavenger 0.015 M HOAc/OAc' 0.015 M HSOL/SOZM, 0.015 M HCOC none 0.1 M HC03-, 0.1 M HOAc, none
___
PH 3.3-8.0 3.0, 5.0 7.0, 8.2 5.0, 6.2 a
pH was not adjusted during these experiments. 0 = OXYGEN OG : OZONE GENERATOR V: VENT T K : T H E R M A L OZONE KILL UNIT F =ROTAMETER
D :OZONE DETECTOR R:REFERENCE SIDE S : SAMPLE SIDE RV: ROTARY VALVE CSTR: STIRRED TANK REACTOR S :SAMPLE POINT
'
pH 2 . 8 . 0.15M H O A C ONE G8T5 LAM? ( X:254nrnl
1 \ 0
Flgure 1. Laboratory ozonatlon system.
E
ance during ozonation in larger reactors of this type. Operating conditions were chosen for experimental convenience and do not represent optimum conditions for actual water treatment. These conditions are summarized in Table I. Ozone in the aqueous phase was analyzed by the indigo method of Bader and Hoigne (19,ZO),using the disulfonate rather than the trisulfonate as originally described by those authors. This method was calibrated in purified water against the iodimetric method of Flamm (21) and checked by UV absorbance with the extinction coefficient of Hart et al. (22). The iodimetric method was, itself, calibrated by quantitative iodine liberation with excess iodide and standard iodate solution, prepared with dried iodate as a primary standard. Ozone in the gas phase was measured by UV absorbance, calibrated against the wet methods by absorbing gas directly into the potassium iodide reagent inside the reactor. Hydrogen peroxide was measured colormetrically by complexation with Ti(1V) (23)or by the method of Masschelein et al. (24). As ozone appears to interfere negatively with hydrogen peroxide measurement using the titanium method, ozone was quickly and vigorously sparged from solution before peroxide measurements were made. The method of Masschelein et al. (24) was not used on ozone-containing solutions. Total oxidants were measured iodimetrically by the method of Flamm (21) but with the addition of a small quantity of ammonium molybdate (25) to catalyze the reaction with peroxides. Water was purified by an ion-exchange/carbon filtration system, after which it was distilled from alkaline permanganate. Total organic carbon (TOC) levels of the final product water ranged from 60 to 300 pg/L as measured by a Dohrmann DC-54 low-level carbon analyzer, with 100 pg/L being an average value. With a dedicated operator and frequent maintenance, precision levels of &0.8% were routinely obtained in TOC measurements in the range 1-3 mg/L. Other chemicals were of reagent grade and were used without further purification.
E!
Results Five sets of experiments were run. These sets are summarized in Table 11. The first set of experiments was run in order to reproduce the results of Taube (15) under conditions where the time dependence of the various analyte concentrations could be determined. Shown in Figure 2 are the results of a typical experiment where one G8T5 UV lamp (primarily X = 254 nm) was used. Ozone, hydrogen peroxide, and total oxidants were measured as a function of the length of time ozone was bubbled into an irradiated aqueous solution of 0.15 M acetic acid. In 762
Environ. Sci. Technol., Vol. 22, No. 7, 1988
15-
tU
I I :
t-
z
W
10-
u
z
0 V
t-
z a
5-
T I M E , min
Flgure 2. Ozone photolysis in acetic acid solution. Curve A is hydrogen peroxide as measured by the method of ref 27. Curve B is total oxidants minus ozone. Curve C is ozone as measured by the indigo method. Curve B is reported as hydrogen peroxide.
0 pH
:
3.3
0 p H : 5.0
0 p H : 6.0 0 p H ' 8.0
-1 \
15
0
E
10
5
0
0
10
20
30
40
50
T I M E , rnin
Figure 3. Effect of pH on hydrogen peroxide accumulation. Total HOAcIOAc- concentration is 0.015 M.
these experiments, hydrogen peroxide (asmeasured by the titanium method) accumulated linearly (curve A), while the ozone concentration remained at about 1mg/L (curve C). The line labeled B represents total oxidants (as hydrogen peroxide) after subtraction of the ozone concentration, as measured by the indigo method. The difference between total oxidants and hydrogen peroxide varies between 0.5 and 0.6 mg/L and is probably within experimental error, since the uncertainties in three different analyses (actually four measurements, since the indigo method involves a bleaching) are involved. The above experiment was repeated in 0.015 M acetic acid with essentially identical results. It was then repeated
0 0.015M H O A c l N o O A c , P H 5
W
pH 7.0
0: [ H p 02],0.001 M HCO;
N
s 4t
0
0 0.015 M H 2 S 0 4 / N a 2 S O q , p H 3
PX
-----------------
0 0 0
10
9
20
n
I n 30
n!n 40
I 50
T I M E,min
T I M E , rnin
Figure 4. Hydrogen peroxide accumulation in the presence of acetic and sulfuric acids.
Figure 5. Ozone photolysis in the presence of bicarbonate.
+
at several pH values, holding the total HOAC OACconcentration constant at 0.015 M. These results are shown in Figure 3, where net hydrogen peroxide accumulation fell off a t near neutral pH values. The maximum hydrogen peroxide accumulation rate calculated from the results in Figure 3 was 0.39 mg/(L.min) or 1.1 X mol/(L.min), in excellent agreement with the value of maximum mass transfer rate in the absence of UV, given in Table I. In order to demonstrate that peroxide accumulation was not simply the result of the stabilizing effect of low pH on ozone and hydrogen peroxide, another set of experiments (series 11)was run. In these experiments acetic acid was replaced by sulfuric acid at pH 3 and 5. The results in Figure 4 demonstrate that a low but measureable peroxide accumulation occurred in sulfate media as compared to that in the acetate runs. After reaching a maximum value, this peroxide decreased very slowly with time. Another series of experiments was run in which acetic acid was replaced by another hydroxyl radical scavenger, bicarbonate ion (series 111). These data are given in Figure 5 where the mass-transfer curve, obtained previously in pH 7 distilled water, is superimposed for reference. As seen in Figure 5, a small amount of peroxide accumulated at the beginning of these experiments, diminishing to the detection limit within a few minutes. Experiments similar to those described above were performed, in distilled water with no added scavenger (series IV). These results are shown in Figure 6 , where the slightly acidic pH values are probably due to absorpton of small quantities of carbon dioxide from the atmosphere. The bicarbonate levels in solution, calculated from these pH values, are too low to account for scavenging by bicarbonate of more than 2% of any hydroxyl radicals formed. The bulk of the hydroxyl radicals should be scavenged by the ozone and hydrogen peroxide concentrations which were present. After an initial rise and fall, both ozone and peroxide concentrations leveled out at easily measurable values. The behavior of the bicarbonate system, described above, led to one final set of experiments (series V), which consisted of hydrogen peroxide photolysis in combinations of the absence and presence of oxygen, acetic acid, and bicarbonate. Since Baxendale and Wilson (26)had shown
2 3
0 L
tU
/*
(L
t-
z 0 0
24 ’ D X
0 0 0
IO
20
30
40
50
T I M E , rnin
Figure 6. Oxidant concentrations during ozone photolysis in distilled water.
that the presence of oxygen had a significant effect on the extent of peroxide regeneration during peroxide photolysis experiments, all of these experiments were run under three different sets of conditions: “oxygen sparged” (oxygen bubbling through the solution during the experiment), “nitrogen sparged” (nitrogen bubbled before and during), and “no sparging” (water from the laboratory distilled water reservoir, with no sparging). As shown in Figure 7, the disappearance curves were first order in all but one of the experiments. Furthermore, the apparent rate constants obtained from the slopes of the first-order plot in Figure 7 are identical within experimental error, for all of the “H20zonly” and Hz02/bicarbonate runs. The slope from the HzOz/aceticacid f nitrogen-sparged experiment is slightly greater than half of that from the above runs, and that from the H,O,/acetic acid/oxygen-sparged experiment is considerably lower still. The data in the early part of the unsparged H202/acetic acid experiment follow those from the oxygen-spargedrun, but later the rate increases until the slope of the log plot Envlron. Scl. Technol., Vol. 22, No. 7, 1988
763
drogen peroxide from which the hydroxyl radical was produced. Thus, peroxide production would be expected in the present system (and in Taube's work) with either eq 1 or eq 2 as the initiation step. Comparison of the hydrogen peroxide photolysis rate in the absence and presence of acetic acid (columns 1 and 3, Table 111)confirms this peroxide regeneration, which shows up as a lower rate of peroxide disappearance. The results shown in Figure 3 are qualitatively consistent with the finding of Staehelin and Hoign6 (27) that increased HOT concentration due to peroxide dissociation (eq 4) at higher pH values results in increased mutual H202
H202 only Nitrogen - Sparged
* HOAc
A a
-1.21 No Sparging I Oxygen - Sparged
0
Y
10
5
15
H 0 2 a H+
20
TIME, minutes
Flgure 7. Firstorder plot of hydrogen peroxlde photolysis data in the presence and absence of radical scavengers.
nitrogen spargad no sparging oxygen sparged
5.7 5.6 5.7
3.0 2.6u 0.86
"Initial HzOz concentration was 1.0 X M. bAll correlation coefficients were in the range of 0.9993-0.9999. 'Limiting value of slope at t > 14 min; see Figure 7.
for those data approaches that of the nitrogen-sparged experiment. The apparent rate constants determined from the data in Figure 7 are given in Table 111. These values are not true rate constants for any one reaction, but represent overall rates for a series of reactions, and are dependent on the particular reactor and UV intensity used. Since these factors were held constant during this series of experiments, it is the relative values of the rate constants which are of interest.
Discussion The results shown in Figures 2 and 3 demonstrate that relatively high concentrations of hydrogen peroxide may be produced in the photolytic ozonation system. This important point has been overlooked in previous studies of photolytic ozonation. Peroxide was produced at a rate which very closely matched the rate of ozone mass transfer into solution, in agreement with the findings of Taube (15). These observations are consistent with either eq 1 or eq 2 as the initiation step (case I or case 11),since Baxendale and Wilson (26) demonstrated that upon photolysis of hydrogen peroxide in acetic acid solution, reaction 3, hyH202-k 2'OH
(3)
droxyl radical reaction with acetic acid in the presence of oxygen, resulted in the regeneration of most of the hy764
Environ. Sci. Technoi., Vol. 22, No. 7, 1988
(4)
03- + H 0 2
(5)
pK = 4.8 f 0.1 (ref 29) (6)
Superoxide reacts very quickly with ozone (30,33)to yield ozonide:
-
03+ O2
(7)
Recent papers by Biihler, Staehelin, and Hoign6 (30)and by Sehested, Holcman, Bjergbakke, and Hart (31-33) have provided values of rate constants for reactions between 'OH and O3 (eq 8), H 0 2 and O3 (eq 9), the forward and reverse reactions in the protonation of ozonide ion (eq lo), and the decomposition of HOBto yield hydroxyl radical (eq 11). Reaction 9 is comparatively slow (32).
apparent rate constant, min-' x lo2 no 0.1 M 0.1 M scavenger bicarbonate acetic acid 5.5 5.3 5.5
-
+ 0202-+ O3
Table 111. Observed Hydrogen Peroxide Photolysis Rate Constants in the Absence and Presence of Scavengersa solution conditions
pK = 11.6 (ref 28)
H02- + O3
@
-1.4
0
+ HO2-
destruction of ozone and hydrogen peroxide. Those authors found that in ozone/peroxide solutions, where even a small amount of the peroxide is dissociated to the anionic form, the dominant initiation reaction for ozone decomposition may be that of ozone and the peroxy anion, which results ultimately in hydroxyl radical and superoxide formation:
HCOi
A
Ht
'OH
-
+ O3
H 0 2 + O2
(8)
+ O3 HO' + 202 Ht + 03- a H 0 3 HOB HO' + O2
HOz
(9)
-
(10) (11)
In addition, Staehelin and Hoigng (34-36), from information in the radiochemical literature, have provided a clear (although in practice, complicated) picture of the radical cycle in the so-called "indirect" ozonation reactions. In these reactions, superoxide is formed from the reaction of hydroxyl radical with many organic compounds in the presence of oxygen (37). Finally, self-reaction of hydroxyl radicals (38) and H02/02-(29) can produce hydrogen peroxide, which can also scavenge hydroxyl radical (39): 2'OH
--
H202
+ HO2' HO2' + 02H202 + 'OH HO2'
-
H202 + 0
(12) 2
+ HOT HO2' + H2O 0 2
(13) (14) (15)
Thus, case I consists of reaction 1followed by reactions 3-8 and 10-15, while case I1 is reaction 2 followed by the same suite of reactions. The entire set of reactions for case I is shown in Scheme I. The simplest 03/UV data set to interpret should be that obtained in the absence of scavenger. However, numerical integration of the rate equations resulting from Scheme I was not possible for two reasons. First, the extent of
first solution is clearly negative. The positive solution can be demonstrated to be too high by calculation of the ozone/superoxide reaction rate, using the experimental M: value of ozone concentration, [O,] z
Scheme I
O3+ HzO k Hz02 + 0 2 H202k 2'OH HzOz O3
G
+ Ht + HO; + H+ 03-+ O2
HOP-
+ HOL
+
H02' F2 0,
O3+ 0,
-
0,
+ H+ ~t HOB HOB 'OH + 02 'OH + HzOz H20 + HOz' 'OH + O3 O2+ HOz' 2'OH H202 2H0; H2Oz + 0 2 HzO + HO; + 02--+ H2Oz + 02 + -OH
0;
+
-
+
-+
-+
rate = R, rate = R, K4 k5
k7[031[02-1= (1.6 x 109 ~-1.~-1)(10-5 ~ ) ( 2 . 4x = 3.8 x 10-4 ~ . s - l = 2.3 X M-min-l
K6 kl
KlO 1215
k8 kl2
k13 k14
where = 10-PH+P&
(17)
The latter solution in eq 16 gives unrealistic values of ['OH] r 7 X lo4M. The first solution in eq 16 gives more reasonable values; e.g., ['OH] < 3 X 10-5/(4 X lo9 X = 7.5 X M, where the inequality is used since the applied ozone dose rate was used as an upper bound for R,, the ozone photolysis rate. This analytical solution further leads to the result
(21)
This reaction rate exceeds the maximum ozone mass transfer rate by 3 orders of magnitude, and thus no ozone residual could be maintained at this superoxide concentration. A similar argument can be applied to the steady-state equation for hydrogen peroxide in case 11:
k1z[OHl2 - ki5Wz021 [OH1 - Rp (22) Reaction between O3 and HOz- is negligible under these conditions. Using kI3 = 8.3 X lo5 M-l-s-l (29), kl4 = 9.7 x lo7 M - W (29), klZ = 4 X lo9 M - W (38), and d[HzOz]/dt = 0, it is easy to demonstrate that unreasonable values of [OH] and/or [OZ-] are required for the two HzOz production terms to equal disappearance by term 3 above. Put another way, case I1 provides no suitable means of regenerating peroxide as quickly as it is used up by reaction with hydroxyl radical. We conclude from the above arguments that photolysis of ozone in aqueous solution yields hydrogen peroxide. Hydroxyl radical is then generated in the secondary reactions represented by eq 3,7,10, and 11 and to a lesser extent by eq 5, 10, and 11. Another interesting result comes out of the above analytical solution, which predicts that both peroxide photolysis and reaction between ozone and HO; decrease to unimportantly small values under the steady-state conditions in Figure 6. The mutual chain decomposition comprised of eq 6-8,10,11, and 15 dominates the reaction system under these conditions. Peroxide photolysis is, however, necessary for initiation of the reaction at the pH values shown in Figure 6, since the 03/H02- reaction is too slow. The lack of accumulation of hydrogen peroxide in the experiments where sulfate replaces acetate (Figure 4) is explained by scavenging of hydroxyl radical to form sulfate radical anion, SO4'- (38): 'OH
where D, is the "utilized" dose rate, Le., that ozone which was removed from the gas stream, and p is the relative rate of *OH scavenging by ozone and hydrogen peroxide. Substituting steady-state ozone and peroxide concentrations from the data at pH 6.2, Figure 6, and rate constants k, = 1.1X 10s M-'*s-' (32) and k15 = 2.7 X lo7 M-'.s-' (39) leads to the reasonable result Ro s 0.30,
M)
kll
ozone and peroxide photolysis at any given instant during the experiment is not known. Second, since a chain reaction is involved, very low concentrations (10-11-10-12M) of active species ('OH and Oz-) are maintained by small differences between relatively fast production and destruction reaction rates, typicaly on the order of lo-' M d . Uncertainties in literature rate constants and the experimental data are large enough to cause instabilities in this system of "stiff" differential equations when numerical values are substituted into them. However, the following kinetic arguments may be made regarding the initiation step in photolytic ozonation. The rate equations corresponding to case I (Scheme I) and case I1 (Scheme I with the first equation replaced by reaction 2) can be solved analytically for the steady-state regions shown in Figure 6. This process involves solution of a quadratic equation for ['OH], which leads to the two solutions
Ly
10-8
(19)
which says that at steady state about one-third of the ozone transferred into the liquid was photolyzed. Similar analytic solution of the rate equations resulting from case I1 leads to the solutions
where P = ks[031/ki5[H2021 and 7 = k7[031/(k15[Hz021 + k,[03]). Numerical evaluation of the second solution gives [02-] 2.4 X lo-, M, which is much too high, while the
-
+ S042-
*S04-l+ -OH
(23)
The fraction of hydroxyl radical which participates in this reaction in, e.g., the experiment at pH 5, Figure 4, is
f =
k23['OH] [so4? ~Z~[*OHI[SO?-I + kd'oH1 [031+ M ' O H I [HzOzI (24)
= 0.95
where the value 1223 = 1.6 X lo6 M-l-s-' (38)has been used. No report of a reaction between ozone and sulfate radical anion was found. An electron-transfer reaction seems unlikely because of the electrophilic nature of ozone. Sulfate radical does, however, react with hydrogen peroxide, with k25 = 1.2 X lo7 M-l-s-' (40):
-
'SO4- + H202
HS04- + H02'
(25)
Thus, sulfate radical destroys peroxide while the superoxide formed from H O i destroys ozone and produces more hydroxyl radical. The results in Figure 5 are explained by a similar mechanism. Reaction of hydroxyl with bicarbonate proEnviron. Sci. Technol., Vol. 22, No. 7, 1988
765
x O2
J RHTHo-HRH
Flgure 8. Reaction cycles in photolytic ozonation.
duces carbonate radical anion, TO;, with kzs = 1.5 X lo7 M - W (41),which then reacts with hydrogen peroxide with k27 = 8 X 10’ M-b-’ (42): ‘OH
+ HC03-
‘COS-
-
H 2 0 + ‘COB-
+ H202
products
(26) (27)
Comparison of the data in column 2, Table 111, with those in column 1 indicate that the product of reaction 27 is capable of regenerating hydrogen peroxide, by analogy with the unscavenged case (reactions 3, 15, and 13 and 14), column 1. Otherwise, the sequence of reactions 3,26, and 27, where the products are “inert”, would lead to a peroxide disappearance rate 1.5 times as great as that in the absence of scavenger. We suggest that the product of this reaction may be H02’: ‘CO,
-
+ H202
HC03- + H02’
(28)
which then regenerates hydrogen peroxide by eq 13 and 14. The difference between the rate constant in the nitrogen-sparged H202/aceticacid experiment and half the value of the rate constants found in the H202/HC03-experiments (column 2) represents the extent to which hydrogen peroxide is not regenerated by the products of hydroxyl radical attack on acetic acid, and amounts to about 5%. This “nearly quantitative” conversion of the hydroxyl radicals back to hydrogen peroxide is the point which makes interpretation of Taube’s conversion data (15) ambiguous. From the above, we conclude that the major mechanistic pathway for the photolytic ozonation system in the absence of additional hydroxyl radical scavengers such as bicarbonate or sulfate is a radical chain reaction which feeds off of incoming ozone via the ozone/superoxide reaction and maintains small but real ozone and peroxide residuals. This cycle, shown schematically in Figure 8 for an organic compound which reacts with hydroxyl radical by hydrogen abstraction, is essentially the same as that given by Staehelin and Hoign6 (35,36)for ozone decomposition in water, with the addition that the reaction is initiated by ozone being driven to hydrogen peroxide in the photolysis step. Under these “initiation” conditions, hydroxyl radical is formed by reaction of O3 with H02- (eq 5 ) and by peroxide photolysis (eq 3). Once the cycle is initiated, it continues by the pathway indicated by the lower perimeter of the circle in Figure 8 provided that the adduct ‘02MH is capable of releasing a proton to facilitate the elimination of 02-.When this part of the cycle is predominant, the hydroxyl radical yield per ozone molecule may approach unity and ozone photolysis may become relatively unimportant. If superoxide cannot be eliminated, the adduct decays by other (generally slower) reactions, including self-reaction (see Swallow, ref 43). Many of these reactions result ultimately in the regeneration of hydrogen peroxide. 766
Environ. Sci. Technol., Vol. 22, No. 7, 1988
Under these conditions, photolysis of hydrogen peroxide may be an important source of hydroxyl radical throughout the reaction. Ozone and/or hydrogen peroxide may accumulate until the “short-circuiting” reactions (eq 8 and 15, not shown in Figure 7) become significant. Under these conditions, the effective yield of hydroxyl radicals, that is, those which eventually react with substrate rather than with ozone or hydrogen peroxide, may be drastically diminished. This is also the case when substrate is at very low concentrations relative to the ozone input rate, as was seen for tetrachloroethylene in paper I of this series (2). When hydroxyl radical scavengers such as bicarbonate and sulfate are present at significant concentrations, the resulting radical anions may attack peroxide preferentially over ozone. If organic compounds are present, the loss in treatment efficiency due to the presence of the scavengers will depend on the relative rates of reaction of radical anion with organic compound compared to that with ozone and/or peroxide (44). It is thus clear that the efficiency of photolytic ozonation for destruction of organic substrates may vary, depending on the relative ozone input, UV input, substrate concentration, and secondary reactions of substrate and other solutes which are present. This is in accord with the widely varying reaction efficiencies reported in the literature, where there is little evidence that consideration of the rate of photon input [in einsteins/(L.min)] relative to the ozone input rate [in mol/(L.min)] has occurred in anything but a strictly empirical manner and where measurement of accumulated hydrogen peroxide has been nonexistent for ozone/UV systems. Because of the complicated nature of the reaction system in photolytic ozonation, an engineering model is needed in order to accurately predict optimum conditions under which to carry out photolytic ozonation. The engineering model should incorporate the kinetic model described above as well as mass transfer, photo absorption, and participation of the reaction products in subsequent active species generation. Such a model is currently under development and verification and will be reported in a later publication. The same model would of course be useful for 03/H202, 03/H202/UV,and H202/UVprocesses by setting appropriate terms to zero, since the present results unify the photolytic ozonation system with that of base-catalyzed ozonation and 03/H202,as elucidated by Staehelin and Hoign6 (27,30,34-36). Thus, an advantage of the photolytic ozonation system for water treatment is that there are multiple pathways for hydroxyl radical generation, so that the reaction system can “adjust” the mechanistic pathway to suit reaction conditions.
Conclusions (1) Hydroxyl radical is the principal active species in photolytic ozonation. (2) Photolysis of aqueous ozone directly yields hydrogen peroxide, which then, along with ozone, participates in secondary reactions to produce hydroxyl radical. (3) In systems where superoxide is regenerated by the reaction of organic radicals with oxygen, the ozone/ superoxide reaction may become the dominant producer of hydroxyl radical. (4)Optimum conditions for photolytic ozonation can vary widely with substrate type and concentration. A comprehensive kinetic model of the system can provide insight into these optimum conditions. (5) Secondary reactions of radical anions formed in scavenging processes may be important in a continuously sparged system which is driven by constant ozone input.
Acknowledgments
The laboratory work was performed by Dean Meldrum, Vicki van Antwerp, Richard Cotten, and Michelle Smith. The cooperation and assistance of the project officer, J. Keith Carswell, is gratefully acknowledged. Registry No. Os, 10028-15-6; HzOz, 7722-84-1.
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Receiued for reuiew August 14,1986. Revised manuscript received July 23,1987. Accepted January 4,1988. This work was funded in part by the Drinking Water Research Division, WERL, U.S. Enuironmental Protection Agency, Cincinnati, OH.
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