54
Ind. Eng. Chem. Fundam. 1904, 2 3 , 54-60
when external mass transfer resistance is negligible. Even with a finite Biot number a sphere and cylinder always have potential geometric instability at certain levels of conversion. Nomenclature a,. b, g, s = stoichiometric coefficients, eq 1 Bz, = k,L/D, = Biot number for mass transfer CA = gas reactant concentration CAb = gas reactant concentration in the bulk CBo= initial solid reactant concentration De = effective diffusivity in solid product layer Da = k s K / D e = Damkohler number K = k,C%= dimensionless inhibition constant K , = inhi ition constant k , = mass transfer coefficient k , = surface reaction rate constant L = half thickness of the solid particle R = position in the solid particle measured from center line r = reaction rate X = solid reactant conversion y = CA/CAB= dimensionless gas concentration Greek Letters y = parameter defined by eq 9 0 = aCB0L/b k,C, = dimensionless time v = shape factor [ = R / L = dimensionless position in the solid particle Subscripts b = bifurcation values
c = at the reaction interface cm = values when y reaches maximum value cM = values at Position of m a x b ” rate
Literature Cited Aris, R. Ind. Eng. Chem. Fundam. 1867, 6 , 315. Beveridge, G. S.G.; Goklle, P. J. Chem. Eng. Sci. 1868, 23, 913. Cannon, K. J.; Denbigh. K. G. Chem. Eng. Sci. 1957a, 6 , 145. Cannon, K. J.; Denbigh, K. G. Chem. Eng. Scl. 1857b, 6 , 155. Claassen, W. A. P.; Bloem, J. J. Crystal Growth 1960, 50,807. DudukoviE, M. P. Chem. Eng. Sci. 1977, 32, 985. DudukoviE, M. P.; Lamba, H. S. Cbem. Eng. Scl. 1878, 33, 471. Erk, H. F.; DudukoviE, M. P. Ind. Eng. Chem. Fundam. 1883, 22, 55. Ishida, M.; Wen, C. Y. Chem. Eng. Sci. 1868, 23, 125. Kurosawa, T.; Hasegawa, R.; Yaglhashi, T. Nippon Klnzoiu Gakkalsbi 1870, 34, 481. Levenspiel, 0.“Chemical Reaction Engineering”, 2nd ed.;Wiley: New York, 1972; p 361. Matsuura, T.; Kato, M. Chem. Eng. Sci. lg67, 22, 171. McKewan, W. M. Trans. Metall. Soc. 1962, 224, 387. Perlmutter, D. D. “Stability of Chemical Reactors”; Prentice-Hall, Englewood Cliffs, NJ, 1972; p 19. Shen, J.; Smith, J. M. Ind. Eng. Chem. Fundam. 1965, 4 , 293. Smfih, J. M. “Chemical Engineering Kinetics”, 3rd ed.; McGraw-Hill: New York, 1981; p 642. Sohn, H. Y.; Szekely, J. Can. J. Chem. Eng. 1872, 5 0 , 674. Sohn, H. Y.; Szekely, J. Chem. Eng. Sci. 1973, 2 8 , 1169. Szekely, J.; Evans, J. W.; Sohn, H. Y. “Gas-Solid Reactions”; Academic Press: New York, 1976; p 89. Walker, P. L., Jr.; Rusinko, F., Jr.; Austin, L. G. Adv. Cafal. 1959, X I , 133. Wen, C. Y.; Wang, S. C. Ind. Eng. Chem. 1970, 6 2 , 30.
Received for review October 29, 1982 Accepted August 4, 1983
Kinetic Behavior of Ozone in Aqueous Solutions of Substituted Phenols Mlrat D. Gurol’ and Seyyedhassan Nekoulnalnl Civil Engineering Department and Environmental Studles Institute, Drexel University, Philadelphia, Pennsylvania 19 104
Reaction rates of ozone with hydroxylated and methylated phenolic compounds in acidic aqueous solution were determined by a dynamic approach which has proved to be a simple and reliable technique for measuring the rates of fast reactions in solution. The results were analyzed to derive the relation between the chemical structure of the substituted benzene ring and its reactivity with ozone. Ozone reacted as an electrophilic agent, and the stoichiometric factor was 1.0.The rate was first order with respect to both ozone and the phenolic compound.
Introduction Wastewaters generated from processes such as oil refining, coke manufacture, and coal conversion contain phenolic compounds in quantities as great as loo00 mg/L (Patterson, 1975). Most of these phenolic compounds consist of phenol and hydroxylated and methylated phenols (Singer et al., 1978). Ozone has been shown to be capable of destroying several phenolic compounds effectively. The literature abounds in information about the process, which has been obtained by combining the reaction rate and the masstransfer rate together in empirical expressions which are system-specific(Eisenhauer, 1968; Gould and Weber, 1976; Anderson, 1976; Yamamoto et al., 1979; Joshi and Shambaugh, 1982). However, for the successful design and operation of ozone reactors in water and wastewater treatment, the kinetic expressions for the ozonation reactions in aqueous solution are needed. 0196-4313/84/1023-0054$01.50/0
Hoigne and his co-workers (1976, 1982) initiated the search for the fundamental rate equations and reported that ozone may react with solutes either by a direct oxidation, in which molecular ozone reacts at electron-rich sites of the solutes, or by an indirect oxidation, whereby hydroxyl radicals resulting from the decomposition of ozone serve as the oxidant in a series of chain reactions. The direct reactions are highly selective in terms of the solutes with which ozone reacts; i.e., the rate constants differ by several orders of magnitudes. For example, ozone reacts with phenol about 650 times faster than it reacts with benzene (Hoigne, 1982). On the other hand, above some critical pH value (usually above pH 5-6), hydroxyl radicals become the predominant oxidants. Although these radicals react very rapidly with organic and inorganic solutes, they exhibit little solute selectivity. For example, the rate constant for phenol is only 3 times greater than the rate constant for benzene (Anbar and Neta, 1967). The 0 1984 American Chemical Society
Ind. Eng. Chem. Fundam., Vol. 23, No. 1, 1984 55
rate of decomposition of ozone and, consequently, the rate of formation of hydroxyl radicals increase with increasing pH (Gurol and Singer, 1982). Therefore, the kinetics and the mechanisms governing the specific solute reaction are markedly dependent upon pH. For many systems encountered in practice, the direct and the indirect reaction pathways may be of comparable importance. By studying ozone consumption in the presence of several relatively slow reacting compounds in batch reactors and at acidic pH values, Hoigne (1982) has established that the rate in the direct reaction of ozone with a solute can usually be expressed by an equation that is first order with respect to the concentrations of both ozone and the solute. A list has been published which contained the reaction rate constant of ozone with several relatively slow reacting compounds, including that of phenol. Our study was performed to add to this limited pool of information. The reaction rate constants of the following relatively fast reactions with ozone were determined: hydroxylated phenols (catechol, resorcinol, and hydroquinone), cresols (0-, m-,and p-methylphenol), and xylenols (2,3-, 2,6-, 2,4-, and 3,4-dimethylphenol). A dynamic approach which simulates the real ozonation systems was applied to determine the relative rate constants. The results were analyzed to derive a relation between the chemical structure of the substituted benzene ring and its reactivity with ozone. Experimental Approach Gurol and Singer (1983) had previously reported that the reaction of ozone with phenol could be called a “fast reaction”, since the process rate under the operating conditions in most experimental or full-scale ozone contactors would be limited by the rate of mass transfer. Hydroxylated and methylated phenols are expeded to be even more reactive toward the ozone molecule than phenol itself, due to the electron donating properties of the extra hydroxyl and methyl groups on the benzene ring. The kinetic expression for such fast reactions cannot be determined by using the conventional batch method in which either the ozone consumption or the solute removal needs to be measured by taking samples from the reaction mixture at various time intervals. Fast reaction kinetics can be studied by commercially available instruments, such as a “Stopped-Flow Kinetic Analyzer.” In such instruments, the changes in the concentrations of reactants and/or products are usually detected by spectrophotometric techniques. However, for the reaction mixture under consideration here, use of spectrophotometric techniques is not appropriate without preliminary separation of the compounds because of the overlappingof the absorption bands of ozone, the phenolic compounds, and the several products of the reaction in the UV wavelength range of 240-280 nm. Hence, a dynamic approach has been adopted to study the reaction rates of the phenolic compounds mentioned above. Ozone gas was continuously bubbled into an ozone contactor which contained aqueous solutions of the phenolic compounds. Under these conditions, the rate of change in the ozone concentration can be expressed as
where square brackets denote concentrations. The first and the second terms represent, respectively, the rates of mass transfer and of the second-order decomposition of ozone (Gurol and Singer, 1982). The last term is the total consumption rate of ozone due to its direct reaction with the solutes. The rate of disappearance of solute i will be
a summation of the rates of the direct reaction by ozone and the indirect reaction by the hydroxyl radical, OH., as follows
In acidic solution, the rate of decomposition and, consequently, the rate of hydroxyl radical formation are negligible (Gurol and Singer, 1982). Therefore, eq 2 can be reduced to (3) Since the pKA values of phenol and of the methylated and hydroxylated phenols are greater than 9, the phenols will be present in solution in their molecular forms. Acidic solutions of pairs of phenolic compounds, or mixtures containing as many as six compounds, were ozonated simultaneously in the ozone contactor. According to eq 3, the ratio of the rates of removal of the compounds would be d[S11 /dt = -kl[SlI d[S21 /dt k2[S21
(4)
or, upon integration between time = 0 and time = t
[Sll
[S2l (5) [SI10 k, [Szlo The apparent ratio of the rate constants was obtained by plotting In ([Sl]/[Sll0) against In ([S2]/[S2l0),according to eq 5. Mass Transfer of Ozone. Equations 1 and 2 were written by assuming that the oxidation reactions will not be fast enough to take place completely in the diffusive f h - t h e liquid region near the gas-liquid interface where mass transfer of dissolved ozone occurs by molecular diffusion. To check this assumption, the kLa, a, and the kL values were determined for the bubble column used in this study. For the details of the experimental setup, refer to Gurol and Singer (1982) and Singer and Gurol (1981). According to Danckwerts (1970), the conditions to be fulfilled if no appreciable amount of gas reacts in the diffusive film with compound i is ln-=-ln-
kl
(7)
A photographic technique used by Zieminski et al. (1967) was adopted for determination of the bubble size inside the column. The pictures of the bubbles formed at a gas flow rate of 0.65 L/min showed an average bubble diameter, d,, of 0.14 cm for the ionic strength of 0.1. The fractional gas holdup, h, was determined by two independent methods. First, the correlation between the gas holdup and the superficial velocity, Vsr, developed by Hughmark (1967) for the air-water system in a 2-in. column was used. From this correlation, h was determined to be 0.027 for a superficial velocity of 0.54 cm/s. Secondly, the liquid height in the column under quiescent conditions and during ozonation with gas bubbles was measured. The value was calculated as 0.025 according to the equation
56
Ind. Eng. Chem. Fundam., Vol. 23, No. 1, 1984
Both values agree well with the holdup of 0.021 determined by Hill and Spencer (1973) for a V,, value of 0.58 cm/s in an ozone reactor. The mathematical average of the two h values was substituted in the following equation to determine the unit interfacial area as 1.11cm-'.
50
5
40
F
Y
.-
30
I
a
= 6h/dp
(9)
20 al
The kLa value for ozone under these conditions was 130 h-' (refer to Gurol and Singer (1982) for the details of kLa measurement). Hence, kL was calculated to be 0.032 cm/s. Additionally, the following empirical equation, given by Hughmark (1967) was used to calculate the kL value as 0.019 cm/s. NSh
= kLdp/DA =
2
+ 0.0187 [N&O.'* X N sc
0.55
NRe =
X
( dp g / D A
dpUsP
P
) Oo4 '
1 (lo) (11)
Here, Us is the slip velocity which is the ratio of the superficial velocity by the holdup when the liquid velocity is zero. The diffusivity of ozone gas, D o , was calculated to be 1.7 X lo4 cm2/s at 20 OC by using the Wilke-Chang correlation (Wilke and Chang, 1955). Both of the calculated and the estimated values for kL are in agreement with 0.020 and 0.017 cm/s, which were reported by Majumdar et al. (1977) and Hill and Spencer (1973) in similar sparged ozone reactors, except that those reactors were operated in a continuous mode with respect to the liquid as well as the gas flow. Hence, according to eq 6 and 7, for an average k L value of 0.020 cm/s and the initial solute concentrations of 5 X to 5 X lo4 M, the z i k , value should be less than 4 X lo5 or 4 X lo4 L/(mol s), depending upon the initial concentration, in order for the reaction to take place in the liquid bulk. In other words, for those compounds with rate constants below the calculated values, the mass-transfer rate of ozone will not be enhanced due to the chemical reaction. Using the rate constant of 1300 L/(mol s) for phenol, as reported by Hoigne (1982),the relative rate constants fulfilling that condition were calculated to be less than 40-400 times that of phenol, depending upon the initial concentrations. However, this dynamic approach should be applicable for even faster reactions if the reaction of ozone in the diffusive film is pseudo first order; i.e., the kLa value is large enough so that the concentration of the solute i is maintained virtually undepleted, with its bulk concentration, [ S i ] , holding right up to the surface of the gasliquid interface. According to Danckwerts (1970), this situation is possible if M I 2
< 1p 2
(12)
where E'=l+-
ZiDi[$1 D*[AI*
Accordingly, for [O,]*= 0.8 mg/L = 1.67 X M, 2 = 1,according to the observations explained later, and [Si] N 5 X M, and if Diis assumed to be equal to Do3 for an approximate solution, then ki is calculated to be 1.0 x lo7 L/(mol s). In summary, under the experimental conditions, the dynamic approach would still be valid for those compounds with rate constants of 400 to 10000 times that of phenol, since the concentration profile of the solute in the diffusive film will not be significantly reduced com-
0 C
0 0
10 0
0
5
1'0
1'5
i0
25
Time ( m i d
Figure 1. Simultaneous oxidation of phenol and m-cresol by ozone.
pared to the bulk concentration, despite the enhancement in the mass-transfer coefficient of ozone. Experimental Procedure The experimental setup was explained in detail previously (Gurol and Singer, 1982; Singer and Gurol, 1981). Pre-measured amounts of phenolic compounds were dissolved before each experiment in distilled, deionized, and pre-ozonated water. The pH was adjusted to 2.5-3.0 by dilute sulfuric acid. Sodium sulfate, which is inert with respect to ozone decomposition in water (Hoigne and Bader, 1976),was used to adjust the ionic strength of the solution to 0.1. Solutions in the reactor were brought to 20 "C by circulating water at constant temperature through the jacket around the reactor. The voltage to the ozone generator was adjusted to obtain 2 or 5 mg of ozone per liter of the gas mixture of oxygen and ozone. Under these conditions, the saturation concentration of ozone was determined to be 0.8 or 2.0 mg/L, respectively, by using the solubility coefficient of 0.41, as determined by Gurol and Singer (1982). The gas flow rate was varied between 0.5 and 1.0 L/min, so that the phenols with initial concentrations of 5 to 50 mg/L would be completely oxidized over a time period of 5-40 min. During the ozonation of the solutions, samples were taken at appropriate time intervab the reactions were stopped immediately by NazS2O3.The samples were kept in a freezer until the time of analysis. A Varian High-pressure Liquid Chromatograph (HPLC) with a CI8reverse-phase column was used for the separation of the phenolic compounds at ambient temperature. An isochratic method with a solvent mixture of 35% acetonitrile and 0.5% acetic acid in water was applied. The phenols were detected simultaneously by a UV detector at a wavelength of 254 nm and a fluorescence detector with excitation and emission wavelengths of 220 and 313 nm, respectively. The peak retention times in minutes for the compounds at a solvent flow rate of 2 mL/min were as follows: hydroquinone, 3.0; resorcinol, 3.4; catechol, 3.8; phenol, 6.0; m-and p-cresol, 8.2; o-cresol, 9.0; 3,4-xylenol, 11.4; 2,3-xylenol, 12.5; 2,6-xylenol, 13.4; 2,4-xylenol, 14.1. Some phenolic compounds were irradiated by UV in a 4-L quartz reactor which was very similar to the bell jar used by Peyton et al. (1982). The reactor fit snugly into a Rayonet Photochemical chamber equipped with 16 lowpressure mercury lamps. The lamps emitted at 254 nm and were rated at 2.2 W each. The samples collected during the UV-irradiation experiments were analyzed by the HPLC as explained above. Concentrations of ozone in the gas and in solution were measured by the KI and indigo procedures, respectively, as reported earlier by Singer and Gurol (1981). Results and Discussion Relative Reaction Rate Constants. Figure 1 shows the results of a typical experiment in which two phenolic compounds were oxidized simultaneously in the ozone
Ind. Eng. Chem. Fundam., Vol. 23, No. 1, 1984 57
0:1
-0
0:2 -In
0.3
El CPlO
Figure 2. Relative reaction rate constant for m-cresol with respect to phenol. 5
50
-
c
E Y
4 s E
2 40
0 E
-
3 5
30
L
C
C
U
0
g
u
20 10
0
3.0
Y
C
'=2
Time ( m i d
Figure 4. Ozonation of phenol and methylated phenols simultaneously.
1
2.0
2 5 Hydroquinone
0
\
PHz2.5
T=20°C
v=o.1
t
log
"i kP 1.o
-
2
4
6
8 1 0 1 2
Time ( m i d
Figure 3. Ozonation of phenol and hydroxylated phenols simultaneously.
0
- 0.4 contactor at pH 2.5. m-Cresol reacted faster than phenol, as would be expected. In this, and all other experiments, the concentration of ozone in the solution remained below the detection limit of the indigo method (Bader and Hoigne, 1981) until the phenolic compounds were oxidized completely. This observation indicated that mass-transfer-limited conditions would prevail during the ozonation of compounds with reaction rates in the same order of magnitude or faster than that of phenol. In Figure 2, the data presented in Figure 1were plotted according to eq 5, which yielded a relative rate constant for m-cresol of 4.4 f 0.8 with respect to phenol, within the 95% confidence interval. The same procedure was repeated for various pairs as well as the mixtures of the compounds. In Figure 3, the concentration profiles for phenol and the p - , 0-,and mhydroxylated phenols are presented as a function of the ozonation time. Due to the higher electron density in the aromatic ring,the hydroxylated phenols reacted faster than phenol. The reactivities of the CH3-substitutedphenols were in the expected order (Figure 4). The dimethylphenols, which reacted fastest, were followed by the monomethylated phenols and finally by phenol. As is evident from Figures 3 and 4, during the simultaneous ozonation of the solutes, the removal rates increased after the fastest reacting solute was oxidized to an extent to allow more ozone to be available for the reactions of the remaining solutes. When the ozonation process was continued further, this behavior was repeated for the subsequent fastest reacting compound. Nevertheless, this change in the slopes of the concentration profiles did not affect the measured relative rates significantly. This is anticipated, since the relative rates are independent of the ozone concentration, according to eq 5.
I
-0.3
-0.2
-0.1
o
a
Figure 5. Relative rate constants for direct ozone reaction vs. Hammett substituent constants.
Table I. "he Relative Oxidation Rate Constants of Substituted Phenols with Respect to Phenol ~
compound
ki/kp
o-cresol m-cresol p-cresol 2,6-xylenol 2,3-xylenol resorcinol 3,4-xylenol 2,4-xylenol catechol hy droquine
4.4 4.4 11 15 19 70 76 76 220 1100
The relative oxidation rate constants were determined by taking the arithmetic averages of a minimum of two measurements made in different reaction mixtures. These relative rate constants are listed in Table I with respect to the rate constant of phenol, k,. As would be expected from direct reactions of the ozone molecule, the constants in Table I differ by a factor of as much as 1100. This observation, which indicates high selectivity for the solutes, contradicts the nonselective reactions by hydroxyl radicals. For example, Anbar and Neta (1967) have reported that the reaction rates of both hydroquine and p-cresol with hydroxyl radicals are only 1.2 times faster than the reaction rate of phenol with the radicals. Hence, our observations prove that in acidic media, ozone molecule is the predominant oxidant for the phenolic compounds. In Figure 5, the relative rate constants are plotted vs., the "Hammett substituent constant", u, based on u(pheno1) = 0 (Klumpp, 1982; Perrin et al., 1981). A "susceptibility factor", po,, of -8.0 f 1.0 was calculated for the compounds
58
Ind. Eng. Chem. Fundam., Vol. 23, No. 1, 1984
included in this study. This reasonable fit of the data to a Hammett plot indicates that at low pH ozone molecule reacts with substituted phenols primarily by an electrophilic attack. Additionally, the hypothesis of significant free-radical reactions in acidic media is rejected again,since the hydroxyl radicals do not seem to obey the rules of an electrophilic attack. The rate constants for the reactions of the hydroxyl radicals with a few aromatic compounds are listed below in order to clarify the argument: The kbenzene= 3.2 x 109, ktoluene= 3.0 x 109, ,k = 3.1 x 109, and kehlorobenzene = 4.2 X lo9 L/(mol s) (Anbar and Neta, 1967; Sehested et al., 1975). A comparison of the relative rate constants in Table I for o- and m-cresols, for 2,6- and 2,3-xylenols, and for 3,4and 2,4-xylenolsleads to the conclusion that the ortho- and meta-substitution by the CH3- group produce the same effects on the reaction rate. For the specific reaction under study, the resonance effects seem to counterbalance the steric effects at the ortho position. Accordingly, the ortho-substituted compounds were plotted in Figure 5, based on the u value reported for meta substitution (note that the u values for ortho substitution are not listed due to the unpredicitibility of the steric effects). As for the hydroxylated phenols, hydroquinone and catechol react faster than resorcinol since the resonance effects activate the benzene ring specifically at the ortho and para positions. However, the reaction is slower for catechol than for hydroquinone, probably because of the steric effects created at the ortho position by the OH substituent. Catechol was not included in the Hammett plot since the u value is not available for the ortho substitution. The inductive effects work in the opposite direction to the resonance effects for OH substitution and, therefore, for resorcinol, a rate constant lower than that of phenol is predicted by the Hammett equation (Klump, 1982). Because of the disagreement between the observed and the predicted rate constant for resorcinol, this compound was not plotted in Figure 5. Bailey (1972) suggested that catechol is produced as an intermediate product of the electrophilic attack of phenol by ozone. Both catechol and hydroquinone were identified and quantified by Singer and Gurol (1981) in concentrations as much as 30 mg/L, in the reaction mixtures of phenol and ozone at neutral and high pH values. In acidic solutions, however, the concentrations of catechol and hydroquinone were less than 1 mg/L. Now, the new findings that catechol and hydroquinone are, respectively, 220 and 1100 times more reactive than phenol with ozone might explain why the concentrations of these intermediates cannot reach significant levels during ozonation of phenol in acidic solutions. Studies with a Radical Scavenger and UV Radiation. The studies explained below were performed for further analysis of the direct reaction pathway observed in acidic media. Hydroxyl radicals react with various compounds, e.g., tert-butyl alcohol, TBA, with rate constants of about lo8 L/(mol s), which is comparable to the rate constants for phenols. However, TBA is almost inert against molecular ozone. Hence, TBA was used in excess amounts as a “scavenger” compound for the hydroxyl radicals. It was assumed that if the free-radical reaction has a significant contribution to the overall reaction mechanism at low pH, the addition of TBA should increase the relative rate constants of the solutes, since the ozone molecule would react with the solutes very selectively. The relative reaction rate constants were determined in the presence of 2.7 x to 0.27 M TBA and were compared to those mea-
Table 11. Relative Reaction Rate Constants in the Presence and Absence of TBA [TBA] = [TBA] = 0
k,lk* hresorcinol
3.7
i
20 000 mg/L 3.5
0.22
* 0.27
k2.3-xylenol k2,3-xulenol k2,6-xylenol k2,6-xulenol
1 . 3 i 0.08
1.4 f 0.16
1.3 ? 0.10
1.1 i 0.06
2 . 5 * 0.19
2.3
kp-cresol kp-cre sol
0.56
i
ko-cresol
.5
4
1
I
25
50
I
75
100
125
Time (min)
Figure 6. Oxidation of phenol and hydroxylated phenols by UV radiation.
sured in the absence of TBA. The amount of TBA to be used in each experiment was calculated to provide a “scavenging power” which is at least ten times the scavenging power of the phenolic compounds. In this study, the product of the rate constant and the concentration of the solutes reacting with the radical is termed the “scavenging power”. The results for a mixture of resorcinol, 2,3- and 2,6-xylenols, and p - and o-cresols, each with an initial concentration of 50 mg/L, are compared in Table 11. For this and most other experiments, the addition of TBA did not affect the measured relative rates significantly within the 95% confidence interval. However, when the pairs of hydroquinone/phenol, catechol/phenol, and resorcinol/phenol were ozonated without TBA, the relative rate constants for hydroquinone, catechol, and resorcinol were measured lower than those listed in Table I. Consequently, the addition of TBA to these solutions increased the value of the relative rate constants to match those in Table I. Further studies are being conducted in our laboratories to understand the reaction mechanisms for these three combinations of compounds. Several different methods of generating hydroxyl radicals in aqueous solution have been reported (Dorfman and Adams, 1973). Among these are the radiation methods in which water is decomposed by electrons, the photochemical methods in which the radicals are formed from water by the absorption of light in the UV range, the chemical methods such as the Fenton’s reaction, and the reaction of titanous oxide with hydrogen peroxide. In this study, different combinations of the phenols were exposed to UV radiation; the concentration profiles of the compounds were compared to those obtained during ozonation. Figure 6 shows an example for a mixture of phenol and hydroxylated phenols. It becomes obvious from the comparison of Figure 6 with Figure 3 that the predominant
Ind. Eng. 7
Values measured in lhis study
OH
OH
OH
e Values reported by Hoigne.1982
2
.I I 0 -0.8 -0.7
cii
.A
"
-6.6 -0.5 -0.4 -0.3-0.2 -0.1
0
0
Figure 7. Absolute rate constants for direct ozone reaction versus Hammett substituent constants.
reaction mechanisms are different in those two cases. For example, phenol is oxidized faster than catechol and resorcinol in the irradiated reaction mixture. On the other hand, some preliminary studies in basic media produced results which are similar to the results of the UV irradiation in terms of the order of the reactivities of the compounds. Further studies with other radical generating methods are now being conducted in our laboratories in order to understand the reaction mechanisms of ozone at higher pH values. Absolute Reaction Rate Constants. A mass balance for ozone and the phenolic compounds around the ozone reactor for several experiments indicated a stoichiometric factor of 1.0 for the reactions; Le., one mole of ozone was consumed per mole of the phenols that were oxidized. As the intermediate products accumulated in the solution and started to consume additional ozone, the stoichiometric factor increased to 3, as also determined by others (Eisenhauer, 1968; Bailey, 1972; Hoigne and Bader, 1976). However, if the solutes were highly reactive with ozone, these reactants were preferred by ozone over the intermediate products. For example, during the first 7 min of the experimental run depicted in Figure 3, a total of 7.0 X lo4 M of phenols (4.5 X M hydroquine, 1.7 X lo4 M catechol, and 0.8 X M resorcinol) were oxidized, consuming exactly 7 X IO4 mol of ozone per liter of solution. This observation is consistent with the hypothesis by Bailey (1972) that the initial attack of phenol by one mole of ozone either breaks the benzene ring or produces a hydroxylated produce, yielding in both cases a different compound than the parent phenol molecule. The above hypothesis was also confirmed by Gurol and Singer (1983) by mathematical simulation of the ozonation system for phenol. Combining the stoichiometric information presented above with the reaction rate constant of ozone with phenol, which is 1300 L/(mol s), (Hoigne, 1982), the absolute reaction rate constants were calculated for the phenolic compounds. These rate constants are plotted together with those of benzene and substituted benzene compounds (Hoigne, 1982) in a Hammett Plot in Figure 7. The data points, which were determined by two different approaches and experimental methods, were best described with a straight line which had a slope of -8.2 F 0.8, which is in agreement with the poQ value presented in Figure 5. The agreement between the dynamic approach presented here and the batch method of Hoigne (1982), in which the reaction rates were determined by measuring the ozone consumption under pseudo-first-order condi-
Chem. Fundam., Vol. 23, No. 1,
1984
59
tions, confirms the observations that the reaction is first order with respect to both ozone and the solute, and that the stoichiometric ratio is 1.0 in these reactions. Additionally, the plot in Figure 7, by covering a range of six orders of magnitude for the rate constants, supports our contention that ozone behaves as an electrophilic agent during its reaction at low pH with the methyl- and hydroxyl-substituted aromatic ring. The dynamic method presented here has proved to be a simple and reliable technique for measuring the reaction rate constants of the phenolic compounds with ozone. Summary and Conclusions The reaction kinetics of various substituted phenols with molecular ozone in aqueous solutions were studied in a semi-batch reactor into which ozone gas was supplied continuously. In acidic solutions, molecular ozone reacted as an electrophile via direct reaction pathway; the contribution of the free-radical reaction to the overall reaction rate was negligible. The methyl and the hydroxyl substituents on the benzene ring increased the rate of ozone attack significantly. The Hammett relation adequately described the substituent effect on the reactivity of the molecules, except for resorcinol. The observations indicated that the reactions were first order with respect to the solute and the ozone molecule, and that the stoichiometric factor was 1.0. The dynamic approach used in this study has proven to be a practical and reliable method for determining the rate constants of the fast-reacting compounds with ozone, with the provisions regarding the mass-transfer characteristics of the ozonation system. In the neutral pH values of most natural waters and wastewaters, the direct ozone reactions and the free-radical reactions might proceed simultaneously at comparable rates. Since the species present in water can catalyze or inhibit the rate of decomposition of ozone, the chemical composition of the waters, e.g., the amount of carbonate alkalinity, will affect the degree of contribution of each reaction pathway. The results of this study, if combined with the kinetic information about the free-radical reactions, can help one to understand the overall reaction mechanism of ozonation, leading to the optimum utilization of this technique for the treatment of water and wastewaters. Acknowledgment This study has been partially supported by NSF Grant No. CEE-8204922 and by the Graduate School of Drexel University. We thank Richard E. Speece, Irwin H. Suffet, and William S. Waddell, Jr., for reviewing the drafts of the manuscript. Nomenclature a = interfacial area per unit volume, cm-l [A]* = saturation concentration of the gas, M DA = diffusivity of the gas in the liquid, cm2/s Di = diffusivity of the solute in the liquid, cm2/s Do3 = diffusivity of ozone in water, cm2/s d, = average bubble diameter, cm g = acceleration of gravity, cm/sz h = gas holdup, fractional k A = reaction rate constant of the gas, L/(mol s) k , = decomposition rate constant of ozone in water, L/(mol 8)
ki = reaction rate constant of solute i with ozone, L/(mol s) k i = reaction rate constant of solute i with hydroxyl radical, L/(mol s) kL = liquid-side mass-transfer coefficient, cm/s kLa = liquid-side overall mass-transfer coefficient, s-l k , = reaction rate constant of phenol with ozone, L/(mol s)
Ind. Eng. Chem. Fundam. 1904, 23, 60-64
60
L, = liquid height in the column under quiescent conditions, cm L , = liquid height in the column during ozonation, cm NRe = Reynolds number Nsc = Schmidt number N S h = Sherwood number [O,] = concentration of dissolved ozone in water, M [O,]* = saturation concentration of ozone in water, M [OH.] = concentration of hydroxyl radicals in water, M [Si]= concentration of solute i in water, M t = time, s Us = slip velocity (the ratio of the superficial velocity by the gas holdup when liquid velocity is zero), cm/s V = superficial velocity, cm/s Zz6%stoichiometric factor (number of moles of ozone consumed per mole of solute oxidized)
Dorfman, L. M.; Adams, G. E. "Reactivity of the Hydroxyl Radical in Aqueous Solution"; National Standard Reference Data System, US. Department of Commerce, NS-RDS-NBS-43, 1973. Eisenhauer, H. R. J. Water Po/lut. Control Fed. 1988, 40(11), 1887. Gould, J. P.; Weber, J., Jr. J. Water Po//&. Control Fed. 1976, 48(1), 47. Guroi, M. D.; Slnger, P. C. Environ. Sci. Techno/. 1982, 76, 377. Gurol, M. D.; Singer, P. C. Water Res. 1983, 77, 1173. Hili, A. G.; Spencer, H. T. "Proceedings, First Internatlonai Symposium on Ozone for Water and Wastewater Treatment"; International Ozone Institute, Washington, DC, 1973; p 367. Hoigne, J.; Bader, H. Water Res. 1978. 10, 377. Hoigne, J. "Handbook of Ozone Technology and Applications"; Rice, R. G.; Netzer, A., Ed.; Ann Arbor Science: Ann Arbor, MI, 1982; Vol. 1. Chapter 12. Hughmark, G. A. Ind. Eng. Chem. Prod. Res. Dev. 1987, 6, 218. Joshi, M. G.; Shambaugh, R. L. Water Res. 1982, 76,933. Klumpp, G. W. "Reactivity in Organic Chemistry", Wiiey-Interscience: New York, 1982; Chapter 3. Majumdar, S. B.; Ceckler, W. H.; Sproul, 0. J. AIChE Symp. Ser. 1977, 73, 166.
Patterson, J. W. "Wastewater Treatment Technology", Ann Arbor Science: Ann Arbor, MI, 1975; Chapter 18. Perrin, D. D.; Dempsey, B.; Serjeant, E. P. "pKa Predlctiin for Organic Acids and Bases"; Chapman and Hall: New York, 1981; p 110. Peyton, G. R.; Huang, R. Y.; Burleson, J. L.: Glaze, W. H. Environ. Sci. Technol. 1982, 76, 440. Sehested, K.; Corfitzen, H.; Christensen, H. C.; Hart, E. J. J. Phys. Chem. 1975, 79, 310. Singer, P. C.; Gurol, M. D. "Wasser Berlin '81"; Colloquium Verlag O.H. Hess: Berlin, 1981. Singer, P. c.; Pfaender, F. K.; Chincilli, J.; Maclrowski, A. F.; Lamb, J. C., 111; Goodman, R. "Assessment of Coal Conversion Wastewaters: Characterization and Preliminary Biotreatability"; Report No. €PA-60017-181, U.S. EPA, Washington DC, 1978. Wiike, C. R.; Chang, P. AIChE J. 1955, 7, 264. Yamamoto, Y.; Niki, E., Shiokawa, H.; Kamiya, Y. J. Org. Chem. 1979, 44, 13. 2137. Ziemlnski, S. A.; Carbon, M. M.; Blackmore, R. B. Ind. Eng. Chem. Fundam. 1967, 6, 233.
Greek Letters = viscosity of water, g/(cm s) CT = Hammett substituent constant, dimensionless po = Hammett susceptibilityfactor for reactions of ozone with benzene ring, dimensionless p = density of water, g/cm3 Registry No. o-Cresol,95-48-7;m-cresol, 108-39-4;p-cresol, 106-44-5;2,6-xylenol,576-26-1;2,3-xylenol, 526-75-0;resorcinol, 108-46-3; 3,4-xylenol, 95-65-8; 2,4-xylenol, 105-67-9; catechol, 120-80-9;hydroquinone, 123-31-9;phenol, 108-95-2. p
L i t e r a t u r e Cited Anbar, M.; Neta, P. Int. J . Appl. Radiat Isot. 1987, 76,493. Anderson, G. L. AIChESymp. Ser. 1978, 73, 265. Bader, H.; Holgne, J. Water Res. 1981, 75, 449. Bailey, P. S. "Ozone in Water and Wastewater Treatment"; Evans, F. L., Ed.; Ann Arbor Science: Ann Arbor, MI, 1972; Chapter 3. Danckwerts, P. V. "Gas-Liquid Reactions"; McGraw-Hill: New York. 1970; Chapter 5.
Received for review December 3, 1982 Revised manuscript received September 14, 1983 Accepted October 12, 1983
Chemical Reactions Accompanying Absorption of NO into Aqueous Mixed Solutions of FeII-edta and Na,SO, Elzo Sada," Hldehlro Kumarawa, and Yasuharu Takada Department of Chemical Englneering, Kyoto University, Kyoto, Japan
The absorption of dilute NO into aqueous mixed solutions of Fe"-edta and Na,SO, was carried out by use of a semibatch bubble column at 308 K and 0.10 MPa total pressure. Chemical reactions accompanyin this chemical absor tion were investigated through quantitative determination of gas- and liquid-phase species. Fe' was oxidized to Fe by NO in the resence of SO3'- according to the mechanism 2Fe" (edta) (NO)" : O S 2Fe" (edta)'S03(N0);PFef)" (e&)N O -,; SO3'-. The main reaction product in the liquid phase was HON(S03);4HS034Fe" (edta)*- 2HON(S03);HN ,O , ., Above pH formed by the reaction of 4Fe" (edta)(NO)'7.5, H$J202further decomposed into NO , and H,O. Gaseous NO , was produced from NO coordinated to Fe'I-edta rather than freely dissolved NO. Furthermore, the effect of coexisting SO2 on the degree of removal of NO is discussed.
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Introduction
The process for simultaneous treatment of gaseous NO, and SO, in flue gas streams by use of aqueous solutions of Na2S03with added Fe"-edta chelate, has been developed in several engineering companies and research laboratories. Among complex reactions accompanying this chemical absorption, it is well-known that the reaction of dissolved NO with Fen (edld2-stoichiometrically proceeds to produce nitrosyl complex. The reduction of the nitrosyl complex by HS03-and S032-,however, seems to be rather complicated and different reduction products have been 0196-4313/84/1O23-OO60$01.5OlQ
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reported at similar reaction conditions. Hence, different schemes of the liquid-phase reduction have been proposed (Tanaka et al., 1976; Sato et al., 1980). In practice, it is important to clarify the scheme of the liquid phase reactions in order to establish the procedure for regeneration of used liquor. In the present work, thus, experiments were carried out on absorption of NO diluted with N2 or He into aqueous mixed solutions of Fe"-edta and Na2S03by use of a semibatch bubble column. The scheme of reactions occurring in this chemical absorption system was investigated in 0 1984 American
Chemical Society
