Ind. Eng. Chem. Res. 1989, 28, 1573-1580
1573
KINETICS AND CATALYSIS Advanced Oxidation Processes. Description of a Kinetic Model for the Oxidation of Hazardous Materials in Aqueous Media with Ozone and Hydrogen Peroxide in a Semibatch Reactor William H. Glaze*-1 and Joon-Wun Kang1 Environmental Science & Engineering Program, School of Public Health, University of California, Los Angeles, California 90024 A model is presented that describes the kinetics of the oxidation of micropollutants in water with the combination of ozone and hydrogen peroxide in a sparged, semibatch reactor. The model is based on known reactions of the 03/H202 system plus mass-transfer characteristics of the reactor. The principal kinetic species for micropollutant oxidation is assumed to be the hydroxyl radical. The model is tested and validated in distilled water spiked with an excess of bicarbonate, a known hydroxyl radical scavenger. Essentially all organic compounds are thermodynamically unstable with respect to oxidation, yet condensedphase oxidation processes are not widely used for destruction of organic contaminants in water, wastewaters, or solid wastes. This is due to the fact that oxidation processes are limited by chemical kinetics, with the result that target organics react too slowly with the common oxidants, even those such as ozone that have highly favorable redox potentials. Also, nontarget contaminants in the matrix may consume unacceptable amounts of oxidant; that is, the oxidant demand may be too high to make the process economically feasible. In spite of these limitations, oxidation processes remain an attractive prospect for waste treatment and the subject of intense research. Oxidation processes are particularly attractive in a era when waste reduction is important in that they can destroy hazardous organic contaminants, not simply transfer them to another phase. The ideal process for destruction of hazardous organic wastes would consist of oxidation with dioxygen at ambient temperature and pressure. The Zimpro wet air oxidation process is an example of a process that approaches this ideal (Dietrich et al., 1985), but as yet catalysts have not been developed that will allow the process to be used at room temperature and pressure for oxidation of a broad array of organics. Stable catalysts are needed that will break the oxygen-oxygen multiple bond and produce active oxygen species capable of oxidizing refractory organic compounds in a cost-effective manner. In the absence of such a process, more reactive (and costly) oxidants have been used for waste treatment, including conventional oxidants such as chlorine, ozone, and hydrogen peroxide. Recent research and development work is showing that rate limitations may be retiioved and costs lowered if conventional oxidants are replaced by combinations of oxidants or oxidants with ultraviolet radiation. We refer to these processes as advanced oxidation
Table I. Rate Constants of Some Organic Compounds with OH Radicals” compd M ^ , 10® M'1 s'1 formic acid 0.2 acetaldehyde 2-chloroethanol tetrachloroethylene nitrobenzene
pyridine trichloroethylene chlorobenzene 1-butanol toluene
vinyl chloride benzene 0
Reference: Farhataziz and Ross, 1977.
processes (AOP) (Glaze et al., 1987a,b; Aieta et al., 1988).
Oxygen-based AOPs include ozone with ultraviolet radiozone with hydrogen peroxide, hydrogen peroxide with ultraviolet radiation, and ozone at high pH values. All of these systems owe their enhanced reactivity, at least in part, to the generation of reactive free radicals, the most important of which is the hydroxyl radical. The hydroxyl radical is an extremely powerful oxidant whose rate constants with organic molecules are generally in the range 8"10 M”1 s'1 (Table I). This means AOP treatment of typical organic substrates will be practical, even if the steady-state concentration of OH radicals is only 10'10-1CT12
ation,
M. For example, if we take the case of a substrate M such as tetrachloroethylene whose rate constant with OH radicals is reported to be 2.3 X 109 M'1 s"\ the pseudo-firstorder rate constant for reaction with OH radicals (k0) will be given by k0 = -d In [M]/df = (2.3 X 109)[OH]SS. Accordingly, even if the steady-state concentration of OH radicals is only 10~n M, the value of kQ will be 0.023 s'1 and the half-life of M will be 30 s. The goal of AOP research, therefore, should be to explore ways to form the OH radical at steady-state concentrations in this order of magnitude. Hydroxyl radicals can be formed by the effect of the ionizing radiation on water, and much of our knowledge of the chemistry of OH comes from the radiation research
To whom correspondence should be addressed. Current address: Department of Environmental Sciences and Engineering, School of Public Health, The University of North Carolina at Chapel Hill, Chapel Hill, NC 27599-7400. *
f
0888-5885/89/ 2628-1573$01.50/0
0.5 0.9 2.3 3.2 3.8 4.0 4.5 4.6 6.8 7.1 7.8
©
1989 American Chemical Society
1574
Ind. Eng. Chem. Res., Vol. 28, No. 11, 1989
Scheme I. Principal Reactions in the 03/H202 System ozone
mass
transfer-
chemical reaction
hydrogen peroxide transfer
1
initiation: (1) H202 H02" + H+ (2) HCV + 03 — 03" + H02 (3) OH' + 03 — 02" + H02 (4) H02 ^ H+ + 02-
propagation: (5) 02" + 03 -*
03" + 02 (6) Of + H+- H03 (7) H03 — OH + 02
net equations for OH formation: initiation: 303 + OH' (8b) HOf initiation: 203 + H202 (8a) OH"
11.8 2.8 X 106 M"1 s"1 70 M"1 s"1 p K = 4.8
reference Behar et al., 1970^ Staehelin and Hoigné, 1982 Staehalin and Hoigné, 1982 Staehelin and Hoigné, 1982
1.6 X 109 M"1 s"1 5.2 X 1010 M"1 s"1 1.1 X 106 M"1 s"1
Buhler et al., 1984 Buhler et al., 1984 Buhler et al., 1984
pK
-
-
*
—
(9) OH + PCE -* products (10) OH + HOf - OH- + H02 (11) OH + H202 — H20 + H02 (12) OH + S; -*· products (13) OH + HCOf — H20 + COf (14) OH + C032- — OH’ + COf (15) OH + 03 —* H02 + 02 (16) COf + H202 — H02 + HCOf (17) C03- + Ij - C032- +
=
20H + 402 20H + 302 radical scavenging reactions: Koester et al., 1971 Christensen et al., 1982 Christensen et al., 1982
2.3 X 10s M"1 s-1 7.5 X 109 *1 s"1 2.7 X 107 M"1 s"1 1.5 4.2 2.0 8.0
X 107 M*1 s"1 X 10® M"1 s'1 X 9 "1 s"1 X 105 M"1 s-1
Weeks and Rabani, 1966 Weeks and Rabani, 1966 Buhler et al., 1984 Behar et al., 1970
If
literature (Chutny and Kucera, 1974). It has also been shown that hydrogen peroxide can be a source of OH radicals by reaction with ozone (Forni et al., 1982; Staehelin et al., 1984), through redox reactions initiated by metal ions such as Fe(II) (the Fenton reagent) (Haber and Weiss, 1934) or by direct photolysis of H202 (Baxendale and Wilson, 1957; Nakayama et al., 1979). The photolysis of aqueous ozone which produces H202 in situ (Taube, 1957; Peyton and Glaze, 1986) or the base-catalyzed decomposition of ozone (Staehelin and Hoigné, 1982) is also a OH-radical generator. Among these processes, the direct reaction of ozone and hydrogen peroxide appears to offer the most promise for practical application (Glaze et al., 1987a, b). This paper will describe a kinetic model of the ozone/ hydrogen peroxide process and the results of experiments to test the model in bicarbonate-spiked distilled water. The following paper (Glaze and Kang, 1989) explores the effects of radical scavengers, including those found in natural waters. A separate series of papers explores the byproducts from these oxidation processes and their potential health effects (Glaze et al., 1989).
Kinetic Model for the 03/H202 Process Chemistry of Systems Involving Ozone, Hydrogen Peroxide, and Water. Forni and co-workers (1982) and Staehelin and Hoigné (1982) showed that the conjugate base of H202 can initiate the decomposition of ozone into
hydroxyl radicals. Scheme I shows the sequence of reactions that occur when ozone decomposition is initiated either by HOf or OH". The overall stoichiometry is given in eq 8a and 8b, depending on whether initiation is by hydroxyl ions or hydrogen peroxide. The rate constants for eq 2 and 3 show that initiation by peroxide is much more rapid than with the hydroxide ion when peroxide is present in millimolar concentrations (Staehelin and Hoigné, 1982). Since both processes are proportional to pH, the relative rates of initiation will be pH independent. In a semibatch reactor with ozone entering at a partial pressure of P (atm), the rate of destruction of an organic substrate M may be represented by -d In [M]/df = fc0iM = ks + &m,oh[OH] + &M,os[03] (18)
where ks is the removal rate due to sparging and &m,oh and
the second-order rate constants for the reaction of substrate with OH radicals and 03, respectively. For the substrate considered in this work, the reactions shown in Scheme I will be assumed to be the only mechanism for substrate removal, i.e., &m,oh[OH] » feM,o3[03] + ks. Moreover, it will be assumed that the rate of initiation by peroxide is larger than the rate of initiation by hydroxyl ions, except in those cases explicitly discussed. For the ozone/peroxide system, if the matrix is nearly pure water, if there is a large concentration of some OHradical scavenger other than ozone, and if the steady-state approximation is assumed to apply to the species Of, Of, H03, and OH, eq 19 can be derived. In eq 19, the nu&m,o3 are
[OH]ss
=
_2fe2(10pH-piC)[H2O2][O3]_ ¿µ,
+ (1
-
SPER)(fc13[HCOf] + fc14[C0321)
merator represents the overall rate of formation of OH radicals (from eq 1, 2, and 4-7). The denominator represents the primary OH scavenging reactions in the substrate and bicarbonate-spiked distilled water system, and k13 and ku are the rate constants for the reaction of bicarbonate and carbonate ions, respectively, with hydroxyl radicals. The selectivity (SPER) in the denominator is a term that is necessary in order to account for the fate of the bicarbonate radical produced in eq 13 and 14. As shown, the product of reaction 16 is the hydroperoxy radical H02 or its conjugate base, the superoxide ion Of. This ion may consume more ozone (eq 5) and eventually another OH radical. SPER is defined as the fraction of COf radicals that reacts with H202 as compared to all other routes: PER
__fe16[CQ3-][H202] MC0f][H202]
+
E*;[COf][I,·]
Rate of Ozone Mass Transfer and Consumption. In semibatch reactor, ozone must be transferred from the gas phase into the liquid phase; hence, mass transfer may be a limiting factor under fast reaction conditions. If ozone enters at a partial pressure of P (atm), the rate of change of ozone concentration should follow a
Ind. Eng. Chem. Res., Vol. 28, No.
d[03)/dt
kha(P/H
=
[03]) Zfe,[S,][03] (21) where feLo is the overall mass-transfer coefficient (s'1), H is the Henry’s law constant for ozone (0.0839 atm m3 mol'1 at 23 °C) (Kosak-Channing and Helz, 1983), and [S,·] refers to the concentration of all species that react with ozone with second-order rate constants k¡. In the presence of hydrogen peroxide, if the water matrix does not contain any constituents that would directly consume ozone, the principal ozone consuming reactions are eq 2, 3, 5, and 15. Using steady-state approximations for radical intermediates, one may derived the following expression for the rate of change of the aqueous ozone concentration:
d[03]/dt
kha(P/H
=
-
-
[03]) 2Ml0pH'píc)[H2O2][O3] (fe10(10pH'pK) + fcn)[H202][OH] 2fe15[03][0H] SPER(fc13[HC03'] + -
-
-
-
-
fei4[C032'])[0H]-3fc3[0H'][03] (22) Rate of Hydrogen Peroxide Transfer and Consumption. Likewise, when hydrogen peroxide is added continuously to a reactor at a feed rate of F (M s'1), the rate of change of the peroxide concentration is given by (23) d[H202]/dt F EMSJ [H202] where S„ refers to a species that consumes hydrogen peroxide. Considering that the peroxide-consuming re=
d[H202]/df
=
F-fc2(10pH'p*)[H2O2][O3]
comes =
([^µ, ][^ 0 · /^))/(^µ, [ ]
+ )e13[HC03'] + + + fe14[C032'] (fe10(10pH'p*) An)[H202]) (27)
Equation 27 stipulates that the rate of substrate oxidation will be proportional to the concentration of ozone in the gas phase and the efficiency of mass transfer (kLa) and inversely proportional to the concentration of scavengers. Equation 27 also implies that too large a dose of peroxide may have an inhibitive effect on the substrate oxidation rate. However, if the peroxide dose is such that residual peroxide concentration is low compared to the bicarbonate-carbonate concentration, then the term (&10(10PH-ptf) + £u)[H202] may be neglected. In region II, residual hydrogen peroxide and ozone levels should be low; however, in region III, hydrogen peroxide concentration will increase with time. The time-dependent peroxide concentration can be expressed by rearrangement and simplification of eq 24, yielding
d[H202]/dt =F-0.5kha(P/H)
+ feu)[H202][0H] Sper(MHC03'] + feu[C032'])[0H] (24)
which may be integrated to yield an expression for the peroxide concentration as a function of time: [H202]t
rates symbolized by Doz and F, respectively (mol L'1 s'1). According to the overall stoichiometry, as shown in eq 8b in Scheme I, the stoichiometric equivalence point would be at Doz = 2F. We shall call this region II. In addition, there are two other regions of interest, i.e., where ozone and hydrogen peroxide are in excess, respectively. The three regions of interest are then region I, Doz > 2F; region II, Doz = 2F (stoichiometric point); and region III, Doz, Dqz < 2F. Kinetics of the 03/H202 System at and beyond the Stoichiometric Point. Regions II and III. When the peroxide/ozone ratio is near the stoichiometric optimum, liquid-phase chemistry becomes fast, as indicated by the rate constants for chain reactions shown in Scheme I. For this reason, the reaction rate is expected to be limited by the ozone-transfer rate into the liquid phase. In this region, ozone concentration in the liquid phase is expected to be at steady state (i.e., d[03]/dt = 0) and at a very low value (i.e., [03] « P/H). Therefore, eq 22 may be solved for the term 2/e2(10pH'pK)[H2O2][O3]:
kLa(P/H)
=
(fe10(10pH'p*) +
feu)[H202][0H]
SPER(fc13[HC03'] + fc14[C032-])[0H]
-
3fc3[OH'][Os] (25)
=
-
exp(-st)}
(29)
where r
F
=
-
0.5jkLa(P/H) + SPER[0H]ss(fc13[HC031 + fci4[C032'])} s
=
0.5(fe10(10pH'pK) +
fcn)[OH]ss
and the hydroxyl radical concentration [OH]ss is stant given by eq 26.
a con-
Kinetics of the 03/H202 System When Doz > F (Region I). In region I, hydrogen peroxide will be consumed as fast as it is added to the reactor and the ozone concentration will build. If the hydrogen peroxide is assumed to be at steady state (i.e., d[H202]/df = 0), then the feed rate of peroxide will be equal to its reaction rate, and eq 24 can be solved for the term fc2(10pH'pK) [H202] [03]. This may then be plugged into eq 19 to yield an expression (corrected for rate at [H202/03] = 0) for the steady-state concentration of OH radicals and the pseudo-first-order rate constant for oxidation of substrate M: [OH]ss
=
2F/(feM0H[M] + (1 + SPER)(fe13[HC03'] +
=
2&µ, ^ / (^µ,
]
fcu[C032'])) (30)
+
(1 + SPER)(fe13[HC03'] + feu[C032'])) (31)
At the lower limit of region I (i.e., at peroxide dose rates
zero), SPER is equal to zero, since peroxide concenSPER may become significant when the peroxide dose increases. In order to calculate the time-dependent concentration of ozone in region I, eq 24 may be rearranged to give an expression for the term fe2(10pH'pK)[H2O2] [03]. This may be substituted into eq 22 to yield, with algebraic manipnear
-
where the term 2/e15[03][0H] has been neglected due to the low concentration of both reactants. Equation 25 may be plugged into eq 19 to yield a modified expression for the steady-state concentration of hydroxyl radicals for the 03/H202 system in regions II and III:
[OH]ss
r/sjl
=
-
In an actual treatment application, ozone and peroxide will be transferred into the reactor liquid simultaneously at
-
-
0.5[OH]ss{(fe10(10pH'pK) + feu)[H202] + SPER(fc13[HC03'] + fc14[C032'])} (28)
-
(fe10(10pH~pK)
2fc2(10pH'p*)[H2O2][O3]
1575
The expression for the pseudo-first-order constant (eq 18), neglecting direct reaction with ozone and sparging, be-
-
actions are eq 2, 10,11, and 16, and when a steady-state approximation is used for C03", the expression for the rate of change of peroxide concentration can be obtained:
11, 1989
([kLa](P/H))/(kM> 0H[M] + fe13[HC03'] + feu[C032'] + (fe10(10pH'p*) + *U)[H202]) (26)
tration is negligible. However,
ulation,
d[03]/dt
=
kha(P/H) -2F +
SPER(fei3[HC03-] + fe14[C032'])[0H]ss (kha + 2fe15[OH]ss + 3A3[OH-])[Oa] (32) -
1576
Ind. Eng. Chem. Res., Vol. 28, No. 11, 1989
Since [OHJss is a constant during a run at a constant dose of peroxide and ozone, eq 32 integrates to (33) [03]t z/w[ 1 exp(-wi)] =
-
where z
=
kha{P/H) -2F
+
SPER(fe13[HC031 + fc14[C032-])[0H]Ss w
=
kLa + 2fc15[OH]sS + 3fe3[OH~]
Ozone Mass Transfer. The volumetric mass-transfer coefficient, kLa, in the sparged reactor was measured under both ozone only and ozone/peroxide reaction conditions, symbolized as kLa and kha*, respectively. For the case of the ozone-only system, the time gradient of ozone concentration in the system can be expressed by eq 34, if the ozone decomposition term in eq 21 can be
d[03]/dt
neglected.
(34) kLa\(P/H) [03]| This will be the case if the pH is controlled near =
-
its neutral value and there is a large (4 mM) concentration of bicarbonate. Equation 34 may be integrated directly
to
In
(P/H)|1
=
[03]t
{(P/H)
[03]t)
-
=
-
expHeLa£)j
-kLat + In (P/H)
(35) (36)
where [03]( is the concentration of ozone in the liquid phase at time t. Thus, plots of In {(P/H) [03],| versus time will yield slopes of -kha. Regardless of the mechanism of ozone consumption by chemical reaction or decomposition, the total amount of ozone transferred (D, mol) in time t is given by -
D
rífeLa|(P/H)-[03]f}Vdt
=
Jo
(37)
where V is the liquid volume of the reactor. For the ozone-only case with no decomposition, one may substitute eq 35 into eq 37 and perform the integration to obtain the following expression for kLa: kLa
=
-[In fl
-
D/((P/H)V)\)/t
(38)
If the peroxide dose is less than the stoichiometric optimum (region I), the residual ozone concentrations will build up as a function of time. Under these conditions, the moles of ozone transferred and the ozone residual may be measured and the ozone residual may be fitted to a polynomial equation of the form [03]t a + bt + ct2. This may be inserted into eq 37 and the integration performed to solve for the apparent mass-transfer coefficient, kLa*, in this region. In the ozone mass-transfer-limited region, [03]t is very small. Thus, eq 37 may be solved directly for the value of feLa*: =
kha*
D*/{(P/H)Vt]
=
(39)
The alpha factor, a, is defined as the ratio of the masstransfer coefficient with liquid-phase chemical reaction, kifl*, to the mass-transfer coefficient without reaction, kLa (Gurol and Nekouinaini, 1985): a
=
kifl*/kE
a
Experimental Section Reactor and Associated Apparatus.
(40)
The 70-L stainless steel, sparged, stirred reactor has been described in a previous paper (Glaze and Kang, 1988). The associated apparatus has been modified and a diagram is shown in Figure 1. The principal improvements have been the
1. Schematic diagram of the apparatus used to study the advanced oxidation processes involving ozone, hydrogen peroxide, and ultraviolet radiation.
Figure
addition of a PCI ozone gas-phase monitor (UV absorption ozone monitor, Model HC-NEMA 12) and the addition of two UNIT Instrument mass flow controllers (UFC 1500A). The total rate of gas flow is 2 L/min (STP), and the flow controllers send precisely (±0.2%) 1 L/min to the ozone monitor and to either of two reactors, making it possible for continuous ozone measurements to be made without affecting the gas flow to the reactor. Materials. Arrowhead distilled water (Arrowhead Water Co., Monterey Park, CA) was delivered in 5-gal polycarbonate bottles and used as received. Tetrachloroethylene (PCE) (laboratory grade, Aldrich Chemical Co., Inc.), sodium bicarbonate (certified ACS grade, Fisher Scientific), sulfuric acid (reagent ACS grade, Fisher Scientific), sodium hydroxide (certified ACS grade, Fisher Scientific), and n-hexane (Optima or GC2, Fisher Scientific) were also used as received. A saturated solution of PCE was prepared by stirring the neat liquid with distilled water overnight. Analytical Methods. Tetrachloroethylene was analyzed by liquid-liquid microextraction with GC/ECD (Glaze and Kang, 1988). The hexane:water ratio was 5:10, and the internal standard was 1,2-dibromopropane at 500 gg/L. After extraction, hexane was transferred with a Pasteur pipet to autoinjector minivials. Linearity of the calibration of the GC/ECD system in the range of interest was demonstrated weekly, and at least a single-point verification was done with each batch of samples. Ozone and hydrogen peroxide in water were analyzed by the indigo trisulfonate method (Bader and Hoigné, 1982) and the horseradish peroxidase fluorescence methods (Lazrus et al., 1985), respectively. Procedure for Kinetic Runs. The procedure is essentially the same as described earlier (Glaze and Kang, 1988). The reactor was filled with distilled water, and then a specific volume of bicarbonate and PCE stock solutions was added to give the desired initial concentrations of each. The target initial concentration of PCE was always 50-70 ppb. The stirrer motor was run at 500 rpm for a few minutes to completely mix the reactor contents. Gas flow was initiated through the ozone generator with valves (Figure 1) in positions such that oxygen containing ozone was passing through the ozone monitor but not into the reactor. Samples were taken from the reactor for each parameter to be measured. The current regulator on the ozone generator was adjusted to give the desired dose rate, and then the valves were changed so as to divert ozonecontaining gas through the reactor. Simultaneously the
Ind. Eng. Chem. Res., Vol. 28, No. 11, 1989
1577
Table II. Mass-Transfer Coefficient, kLa, at Various Ozone Doses
yciUAiuc/
2. Plot of data for the determination of the ozone masstransfer coefficient, kLa, of the semibatch stirred tank reactor: rpm, 500; gas flow rate, 1.0 L min"1; temperature, 23 ± 2 °C.
Figure
metering pump was started with a predetermined setting so as to give the desired peroxide dose rate. The temperature of the reactor contents was 23 ± 2 °C. The pH of the bicarbonate spiked distilled water was adjusted to 7.6-7.S prior to oxidation with a 1 N HC1 solution. Samples for ozone residual analysis were taken in 50-mL volumetric flasks, each containing indigo trisulfonate solution, the amount and concentration of which were predetermined so as to give a large absorbance decrease after reaction with ozone. Samples (3 mL) for hydrogen peroxide analysis were taken in 5-mL vials containing 0.45 mL of horseradish peroxidase solution (8 units/mL) and 0.45 mL of 0.35 M potassium hydrogen phthalate buffer (pH 5.5). Samples for PCE analysis were taken in 40-mL vials containing 1-2 drops of 1 % sodium thiosulfate solution. The vials were filled so as to leave no headspace and sealed with PTFE-lined silicon septa and screw caps. Duplicate samples were taken for PCE analysis at each sampling time.
L
„
ozone
mg/(L min)
(w/w)
(1)
(2)
0.226
0.000 0.100 0.200 0.300 0.400 0.500 0.600 0.800 1.000 1.200 1.400
7.02
6.46
0.000 0.217 0.433 0.715
6.58
0.492
_4
1
’
ozone dose,
(3)
0.000 0.200 0.440 0.704 1.590
ab
Ec
7.80 7.78 7.78 7.78
1.00 1.00 1.08 1.10 1.11 1.15 1.16 1.18 1.17 1.17 1.17
1.00 1.22 1.41 1.50 1.59 1.74 1.78 1.88 1.90 1.90 1.90
1.00 1.66
7.74 7.84
1.00 1.15 1.16 1.17 1.18
7.76 7.80
1.00 1.05 1.16 1.17 1.18
1.00 1.29 1.78 1.85 1.85
6.64 7.15 7.31 7.42 7.65 7.72
6.41
7.65 7.72
1.134 1.00
(4)
6.47
6.86 6.94 7.69
1.71 1.85 1.87
0 (1) = calculated from plot of In (P/H [03]() versus t. (2) = calculated from eq 38. (3) = calculated from eq 37 with polynomial fit of residual ozone data. (4) = calculated from eq 39. b a factor = kLa*/kha (kha taken as 6.63 x 10"4 s"1). c Enhancement factor = D*/D (D* measured for 25 min). -
Results and Discussion Mass-Transfer Characteristics of the Reactor. The volumetric mass-transfer coefficient of the sparged reactor # was measured under both ozone-only (kLa) and ozone/ peroxide reaction conditions (feLo*). For the ozone-only system, the mass-transfer coefficient (kLa) was determined by two different methods. In the first method, typically found in most of the Itierature, kha was determined by plotting In \P/H [03]| versus time for various ozone doses. This method is presented in Figure 2, and values are tabulated in Table II. In the second method, kha was determined by measuring the ozone entering the reactor and in off gas from the reactor (eq 38). The data obtained in this method are also listed in Table II. There is some scatter in the kLa values for the ozone-only system, which may be due to the difficulties associated with measuring a highly reactive substance such as ozone. The average value for kha from all of the measurements is 6.63 X 10"4 s"1 (±0.25). Table II also presents the values of kha* obtained under the various ozone/peroxide oxidant doses. At three fixed ozone dose rates, the hydrogen peroxide dose rate was increased through the ozone mass-transfer-limited region. The apparent value of kha* and the a factor increase as the peroxide dose increases, reaching a limiting value beyond the stoichiometric 03/H202 ratio (0.35). The average value for all of the data in the range 0.4 < [H202/03] give a value for kLa* of 7.76 X 10"4 s-1 (±0.05) (Figure 3). The enhancement factor (E) is defined as the ratio of the gas absorption into the liquid phase with simultaneous liquid-phase chemical reaction to the physical absorption
Hydrogen Peroxide/Ozone
(w/w)
3. Apparent ozone mass-transfer coefficients, feLa*, and a factors in the 03/H202 system: ozone dose, 0.23 mg/(L min); bicarbonate alkalinity, 4 mM; temperature, 23 ± 2 °C.
Figure
-
rate alone. For the batch system, E can be expressed as the ratio of the amount of ozone absorbed into liquid volume is a given time when in the presence of peroxide (D*) to the amount of ozone absorbed in the same time in the absence of peroxide (D). For 25-min contact time, the observed enhancement factor for ozone absorption was in the range 1.80-1.90 in the batch reactor, implying that chemical reaction enhances ozone mass transfer. However, the semibatch reactor used in this work is not appropriate for the precise measurement of mass-transfer enhancements, particularly for a complex system such as [03/ H202]. Further studies to elucidate the mass transfer for the chemical reaction regime of this system are in progress. Test of the 03/H202 Model in Bicarbonate-Spiked Distilled Water. Figure 4 is a semilog plot of data for a set of typical ozone/peroxide runs conducted in bicarbonate-spiked distilled water, showing that the oxidation of PCE is a pseudo-first-order process. Table III is a compilation of pseudo-first-order rate constants (fe0) for these runs in which the hydrogen peroxide and ozone dose rates are varied. The results indicate that addition of peroxide accelerates the oxidation of PCE by a factor of
1578
Ind. Eng. Chem. Res., Vol. 28, No. 11, 1989 _
O
:
Calculated
:
Experimental
Hydrogen Peroxide/Ozone 4. Pseudo-first-order plots for the destruction of PCE by 03/H202 process in bicarbonate-spiked distilled water: ozone dose, 0.23 mg/(L min); bicarbonate alkalinity, 4 mM; pH, 7.6-7.8; temperature, 23 ± 2 °C.
Figure
(w/w)
Figure Pseudo-first-order rate constant, k0, of PCE in distilled water versus peroxide/ozone mass ratio: ozone dose, 0.23 mg/(L min); bicarbonate alkalinity, 4 mM; pH, 7.6-7.S; temperature, 23 ± 5.
2
°C.
Table III. Effect of Hydrogen Peroxide and Ozone Gas Concentration on the Rate of Oxidation of PCE° dose,
mg/(L min)
03 0.226 0.226 0.226 0.226 0.226 0.226 0.226 0.226 0.226 0.226
H202 0.000 0.020 0.042 0.061 0.085 0.150 0.205 0.255 0.341 0.605
H202/03 (w/w) 0.000 0.088 0.185 0.270 0.376 0.663 0.907 1.128 1.509 2.677
3.9 10.3 14.9 18.8 24.1 26.8 28.7 25.7 27.1 20.8
0.492 0.492 0.492 0.492 0.492 0.492 0.492
0.000 0.107 0.213 0.352 0.558 0.713 0.950
0.000 0.217 0.432 0.715 1.134 1.449 1.930
5.7 35.0 56.8 69.6 65.7 56.6 57.1
1.000 1.000 1.000 1.000 1.000 1.000 1.000
0.000 0.220 0.440 0.704 1.590 1.980 2.640
0.000 0.220 0.440 0.704 1.590 1.980 2.640
8.9 65.6 106.3 133.0 127.5 124.8 113.0
k¡¡, 10~4 s"1
0 Water: distilled water. Bicarbonate alkalinity: 4 mM (spiked), pH: 7.6-7.8.
7-15, depending on the ozone dose rate. Figure 5 shows how k0 values change as the hydrogen peroxide dose is changed at a given ozone dose. The shape of the curve indicates that the process is ozone mass transfer limited, since increasing the peroxide dose beyond the stoichiometric optimal dose does not further enhance the removal rate. There is even the suggestion that the rate is decreased past a certain point (H202/03 > 2), which is predicted by the process model. Test of the Kinetic Model in Region III. The kinetic model of the 03/H202 process in region III may be tested in two ways. First, we may use measured values of k0 in this region to compute the value of the rate constant for the reaction between hydroxyl radicals and PCE (&pce,oh)> since all other quantities in eq 27 are known. In the calculation of &pce,oh> eq 27 is simplified by the elimination of the terms in the denominator involving substrate M and peroxide. It can be shown that these terms contribute a maximum of 5% to the total magnitude of the denominator of eq 27 under the conditions used in this work. The
Figure 6. Comparison between measured and predicted [H202] in region III: ozone dose, 0.23 mg/(L min); bicarbonate alkalinity, 4 mM; pH, 7.6-7.8.
Table IV. Estimation of IrM0h from the Measured Ir0 ozone
dose,
P/H,
10"5
mg/(L min)
M
0.23 0.49 1.00
9.6 20.9 42.4
^0,PCE>“ 10"4 s"1
22 56 112
(±3) (±5) (±10)
kpcE,OH>6109
M"1
s"1
(±0.3) 2.2 (±0.2) 2.2 (±0.2) 2.1 (±0.2)
[OHW "12 M 1.1
1.9
2.7 5.3 av
Average value of fcojcE with 0.4 > [H202/03]; corrected for th rate at [H202/03] = 0. 6Calculated from eq 27; kLa* = 7.76 X 10~4 = 83.9 L atm mol"1 s"1; total carbonate alkalinity, 4 mM; pH 7.6; (23 °C). cCalculated from eq 18 with Apce.oh = 2.1 x 109 M"1 s"1. “
results of this exercise are shown in Table IV. The average calculated value of &pce,oh is 2.1 X 109 M"1 s"1, within experimental uncertainty of the literature value, 2.3 X 109 M'1 s"1 (Koester and Asmus, 1971). This represents both an initial verification of the model and also a method for determining the rate constants for other substrates with OH radicals. Further use of this procedure is made in the following paper (Glaze and Kang, 1989). The second test of the kinetic model in region III involves the comparison of residual hydrogen peroxide levels with those values computed from eq 29. As noted in the following section, SPER is expected to be close to unity in region III; i.e., almost all of the bicarbonate radicals are expected to react with peroxide by eq 16. By use of this assumption, Figure 6 shows plots of peroxide residuals for various peroxide doses (0.8-1.4 peroxide/ozone mass ratio) at one ozone dose (0.23 mg/(L min)) along with computed values. It is clear that the agreement between the model and experiment is excellent.
Ind. Eng. Chem. Res., Vol. 28, No. 11, 1989 Table V. Evaluation of Selectivity Term the Fate of the Carbonate Radical
Ms' -1
SPEB
Regarding
[OH]ss,6
h2o2
[H202/03] (w/w)
*0,NET>°
o3
10'4 s'1
10'13
78.47 78.47 78.47 78.47
9.80 20.59 29.66 41.67
0.088 0.186 0.268 0.376
6.4 11.0 14.9 20.2
M
3.05 5.24 7.10 9.62
SpER
0.02 0.24 0.32 0.37
6 = Kg,net = k0 k0 (ozone only); data from cTable III. [OH]ss = x 2.1 M 1. 109 s calculated from SPer ^o,net/^pce,oh> ^pce,oh
a
-
1
eq 30.
1579
Values of k0 in region I calculated from eq 31 are shown one case in Figure 5. The term involving SPER is not included in the predicted line in the figure. The agreement is good at the lower peroxide/ozone ratios, but significant error is apparent near the stoichiometric optimum. Hence, the model is not as accurate in regions I and II as in region III. However, it should be noted that the 03/H202 system will be used primarily at or beyond the stoichiometric H202/03 optimum where the model works well. Further research to elucidate the cause of the errors in region I is in progress.
for
Conclusion
7. Comparison between measured and predicted [03] in region I: ozone dose, 0.23 mg/(L min); peroxide/ozone ratio, 0.1 (w/w); bicarbonate alkalinity, 4 mM; pH, 7.6-7.8.
Figure
Fate of the Carbonate Radical in the 03/H202 System. In bicarbonate-spiked distilled water, the principal OH scavengers are the bicarbonate and carbonate ions, the pK for which is 10.25 (eq 13 and 14). The fate of the carbonate radical in region I is not clear, though disproportionation (eq 41) is one plausible route (Behar et al., 1970). C03'~ + C03' —unknown
(41)
In region III, where hydrogen peroxide builds, eq 16 is presumably an important route. The selectivity factor (SPER) is defined as the fraction of carbonate radicals that proceeds by eq 16 as opposed to all routes, including (41). The factor SPER may be evaluated in region I by solving eq 31 using the data in Table III at [03/H202] ratios less than 0.4. The results are shown in Table V. As expected, the fraction of carbonate radicals that react through eq 16 is increasing with relative peroxide dose, approaching unity in the peroxide-rich region III. The value at the stoichiometric optimum is approximately 0.4. Test of the Kinetic Model in Regions I and II. In region I, ozone levels in the liquid phase should increase with time according to eq 33. The value of SPER is changing in region I as the peroxide dose is increased, but the term involving SPER in the parameter z (eq 33) is negligible (
