4108
Ind. Eng. Chem. Res. 1997, 36, 4108-4113
Kinetics of Oxidation of Ammonia in Solutions Containing Ozone with or without Hydrogen Peroxide Chiang-Hai Kuo,* Fang Yuan, and Donald O. Hill Department of Chemical Engineering, Mississippi State University, Mississippi State, Mississippi 39762
This research investigates the kinetics of traditional ozonation and peroxone oxidation of ammonia in alkaline solutions. The ozonation reaction is governed by the direct oxidation of ammonia with ozone molecules, and the overall kinetics is second order, with first order each in ozone and ammonia. The reaction rate increases slightly with temperature, and at 25 °C, the overall rate constant varies from 12.3 to 27.0 1/Ms as the pH increases from 8 to 10. In the presence of hydrogen peroxide, the peroxone oxidation is controlled mainly by the formation of hydroxyl radical and its subsequent radical reactions. The depletion rate of ozone is first order with respect to both concentrations of ozone and hydrogen peroxide but is nearly independent of the ammonia concentration. The overall rate constant increases from 5860 to 133 000 1/Ms in the pH range of 8-10 at 25 °C, indicating that the rate increases with the hydroxyl ion concentration of the exponent of 0.71. The destruction rate of ammonia depends on concentrations of various species, and for wastewaters containing high concentrations of ammonia, the peroxone oxidation process can be effective and economical to achieve a high efficiency for ozone utilization. Introduction
Kinetic Model
Ammonia is used in the manufacture of fertilizers, refrigerants, and many household cleaning agents and stabilizers. The wide applications result in ammonia contamination in air, surface, and underground waters. Ammonia can be removed from water streams by physical, biological, and chemical methods. Chemical oxidation is effective in the treatment of waters with low concentrations of ammonia, but the well-known oxidation by breakpoint chlorination (Lewis, 1980) may result in the formation of toxic chlorinated compounds in treated waters. Another chemical method, ozonation, is a promising method, and unreacted ozone reverts to oxygen without toxic residue. An advanced oxidation using ozone and hydrogen peroxide mixtures as the combined oxidants (known as peroxone oxidation) also may be effective for ammonia removal from wastewaters. A drastic reduction in ammonia concentration at pH 10.6 following lime clarification was observed by Huibers et al. (1969) in their study of wastewater treatment by ozone. Singer and Zilli (1975) and Ikehata (1975) found that the decomposition rate of ammonia increases with pH in alkaline solutions. In a study of the ozonation of ammonia in solutions of pH ranging from 7 to 8, Hoigne and Bader (1978) yielded a second-order rate constant of 20 1/Ms at 20 °C. Improvement in ammonia oxidation by ozone in the presence of bromide ion was observed by Haag et al. (1984). No literature information is available regarding the kinetics of peroxone oxidation of ammonia. Therefore, this study was undertaken to provide kinetic data for the ozonation of ammonia in the pH range of 8-10 at temperatures ranging from 15 to 35 °C and to investigate the kinetics of ammonia oxidation in the presence of ozone and hydrogen peroxide mixtures. A kinetic model was developed and rate equations derived to assess major mechanisms governing the ozonation and peroxone oxidation of ammonia in basic solutions.
As suggested by results of some recent studies (Staehelin and Hoigne, 1982; Tomiyasu et al., 1985), a kinetic model is formulated here by considering the formation of a hydroperoxide ion as the initial step in the decomposition of ozone in an alkaline solution.
* Author to whom all correspondence should be addressed. E-mail:
[email protected]. S0888-5885(97)00208-X CCC: $14.00
O3 + OH- f HO2- + O2
(1)
In the presence of hydrogen peroxide, a weak acid, its partial dissociation in water results in the production of additional hydroperoxide ions (Taube and Bray, 1940).
H2O2 S HO2- + H+
(2)
The hydroperoxide ions react further with ozone molecules to produce hydroxyl free radical.
O3 + HO2- f OH• + O2- + O2
(3)
Once the hydroxyl radical is formed, it may react with many species in chain reactions (Langlais et al., 1991; Hong et al., 1996). The most important reactions (Staehelin and Hoigne, 1982; Christensen et al., 1982) are
O3 + OH• f HO2• + O2
(4)
H2O2 + OH• f HO2• + H2O
(5)
HO2- + OH• f HO2• + OH-
(6)
HO2• S O2- + H+
(7)
OH- + OH• f H2O + O-
(8)
O3 + O2- f O3- + O2
(9)
O3- + H+ f OH• + O2
(10)
O- + H2O f OH• + OH-
(11)
© 1997 American Chemical Society
Ind. Eng. Chem. Res., Vol. 36, No. 10, 1997 4109
The following equilibrium relationship is assumed valid:
H2O S H+ + OH-
(12)
In the presence of ammonia as a contaminant in the solution, it can react with both the hydroxyl radical and ozone molecules as follows:
NH3 + OH• f Qs
(13)
NH3 + rO3 f Rs
(14)
In the direct oxidation of ammonia in reaction 14, it is assumed that r moles of ozone are required for complete conversion of each mole of ammonia. Many other scavengers, if present in the solution, also may participate in the chain reactions (Glaze et al., 1987; Glaze and Kang, 1989); these additional reactions are neglected in this work. Expressions for depletion rates of ozone and ammonia can be derived using the steady-state approximations for OH•, O-, O2-, and O3-, and the equilibrium relationships given in reactions 2, 7, and 12. The depletion rate of ozone can be written as
-d[O3]/dt ) {k1 + a(1 + b)[H2O2]}[O3][OH-] + rk14[NH3]n[O3] (15) and the oxidation rate of ammonia is
-d[NH3]/dt ) (k13/k4)ab[H2O2][NH3][OH-] + k14[NH3]n[O3] (16) In the above equations, the constant a is defined as
a ) 2k3K2/K12
(17)
and b is a function of various concentrations and rate and equilibrium constants.
b ) 1/[1 + {(k5/k4) + (k6/k4)(K2/K12)[OH-]}[H2O2]/ [O3] + (k13/k4)[NH3]/[O3]] (18) Equations 15 and 16 are expressed in terms of the concentrations of ammonia, oxidants, and hydroxyl ion. On the right hand side of eqs 15 and 16, the second term denotes the direct oxidation rate between ammonia and ozone molecules, and a(1 + b)[H2O2][O3][OH-] represents the rate of indirect oxidation of ammonia by the hydroxyl radical. In the absence of hydrogen peroxide, eqs 15 and 16 can be simplified to
-d[O3]/dt ) {k1[OH-] + rk14[NH3]n}[O3]
which is well recognized in the literature as the selfdecomposition rate of ozone in alkaline solutions (Langlais et al., 1991). Experimental Details Ozone was produced from an Ozoteq Laboratory Ozonator (Model L-50) using extra dry, pure oxygen. The ozone-oxygen mixture was introduced into deionized distilled water to oxidize any contaminants, and then the preozonized water was used to prepare aqueous solutions of dissolved ozone and ammonia. The initial concentration of dissolved ozone was determined by both iodometric titration and a spectrophotometric method. An aqueous solution of ammonia was prepared by adding appropriate quantities of ammonium chloride and sodium hydroxide in the preozonized water. Hydrogen peroxide was also added in the ammonia solution to maintain a desired initial concentration of hydrogen peroxide in a peroxone oxidation experiment. By utilizing a stopped-flow spectrophotometer system (High-Tech Model SF-51), two reactant solutions (one contained dissolved ozone, and another contained ammonia and sodium hydroxide with or without hydrogen peroxide) were mixed rapidly (complete within 0.001 s), and absorbance changes of the mixed solution (at the wavelength of 260 nm) were determined and recorded at various reaction times during a kinetic experiment. With excessive ammonia (and hydrogen peroxide, if present) in the solution, the measured absorbances during a reaction can be converted to the concentrations of ozone, the limiting reactant (Espenson, 1981). The dissociation of solid NH4Cl in the solution resulting in an equilibrium among the ammonium ion (NH4+), un-ionized or free ammonia (NH3), and hydroxyl ion (OH-). The concentration of dissolved ammonia (NH4+ and NH3 combined) was determined by an Ammonia Combination Electrode (manufactured by Corning Co.). The fraction of free ammonia in the solution of a fixed pH can be calculated from
[NH3]/{[NH4+] + [NH3]} ) 10pH-14/{Kb + 10pH-14} (22) Using the ionization (dissociation) constant, Kb ) 1.774 × 10-5 at 25 °C (Weast and Astle, 1978), for example, the fractions of free ammonia (NH3) can be calculated as 0.00552, 0.0526, 0.357, and 0.847 respectively at pH 7, 8, 9, and 10. The preliminary tests in this work showed that ammonium and chlorine ions are not reactive toward ozone or hydrogen peroxide or mixtures of the two oxidants and that free ammonia is not reactive toward hydrogen peroxide alone. Therefore, only free ammonia can be oxidized by ozone or mixtures of ozone and hydrogen peroxide.
(19) Kinetics of Ozonation
and n
-d[NH3]/dt ) k14[NH3] [O3]
(20)
Equations 19 and 20 are expressions for the depletion rates of ozone and ammonia, respectively, by direct oxidation in the traditional ozonation process. Without hydrogen peroxide or any contaminants, eq 19 is reduced further to
-d[O3]/dt ) k1[OH-][O3] ) kd[O3]
(21)
The concentration profiles of ozone during selfdecomposition, ozonation, and peroxone oxidation are illustrated in Figure 1 for three experiments carried out at 25 °C in solutions with pH 9 at a constant initial ozone concentration. By self-decomposition (case 1) alone, the ozone concentration declines slowly in the solution, but the depletion rate is accelerated in the presence of free ammonia with or without hydrogen peroxide (cases 2 and 3). In this figure, the ozone concentration is plotted against the reaction time on a semilogarithmic scale, yielding a straight line for each
4110 Ind. Eng. Chem. Res., Vol. 36, No. 10, 1997 Table 1. Overall Rate Constants ozonation, 1/Ms pH
15 °C
25 °C
35 °C
peroxone oxidation at 25 °C, 1/Ms
8 9 10
5.7 7.4 9.9
12.3 16.7 27.0
30.7 34.6 43.0
5 860 21 100 133 000
Figure 1.
run. This illustrates that the overall kinetics is first order with respect to the ozone concentration for each of the three reactions. The apparent rate constant for an individual run can be obtained from the slope of the straight line or, alternatively, from regression analysis of the concentration data. The apparent rate constants are calculated from the regression analysis; the correlation coefficients are 0.98 or larger for all values reported in this work. For the ozonation reaction with excessive amounts of ammonia in an experiment, eq 19 can be integrated to yield
[O3]/[O3]0 ) exp[-k′t]
(23)
Figure 2.
220 nm). Therefore, the stoichiometric ratio, r, is about 4, as reported earlier by Singer and Zilli (1975). Using r ) 4 and n ) 1, the ozonation rate constant, k14, can be calculated from eq 24 for each experiment. Also, eqs 19 and 20 can be rewritten as
where the apparent rate constant, k′, is defined as
k′ ) kd + rk14[NH3]0n
(24)
The decomposition rate constant, kd, is equal to k1[OH-], as shown in eq 21. For a series of experiments carried out in solutions of a constant pH (ranging from 8 to 10) at 25 °C, the apparent rate constant can be plotted against the initial concentration of ammonia on a logarithmic scale to yield a straight line. This confirms that the ozonation reaction is first order with respect to the concentration of ammonia (n ) 1) and that the overall reaction is second order, as reported by the previous investigators (Singer and Zilli, 1975; Hoigne and Bader, 1978). As shown in eq 24, the decomposition rate constant, kd, and the stoichiometric ratio, r, are needed to calculate the ozonation rate constant, k14. The decomposition rate constant at fixed temperature and pH can be obtained from the literature data (Kuo et al., 1977) or from data obtained in the decomposition experiments carried out in this work, as demonstrated in Figure 1 (case 1). Several tests were conducted to determine the stoichiometric ratio of the ozonation. The results showed that 3.9-4.6 mol of ozone were consumed for conversion of each mole of ammonia and that the reacted ammonia was converted completely to nitrate, NO3- (determined by the spectrophotometric method at the wavelength of
-d[O3]/dt ) {kd + 4k14[NH3]}[O3]
(25)
-d[NH3]/dt ) k14[O3][NH3]
(26)
and
For a series of experiments conducted at identical temperature and pH, the calculated values (k14) were averaged to yield an average constant shown in Table 1. The average deviation of the individual rate constants from the average value at any given pH and temperature is 15%. The results in the table indicate that the ozonation reaction is slow and that the reaction rate increases slightly with pH and temperature. As mentioned earlier, Hoigne and Bader (1978) gave the rate constant of 20 1/Ms for the ozonation at 20 °C in the pH range of 7-8. They also suggested that above pH 9 the oxidation mechanism may be controlled by the hydroxyl radical reaction of ammonia. The gradual change in the ozonation rate with pH (k14 is proportional to [OH-]0.12, as shown in Figure 2) at a given temperature seems to suggest that in the pH range of 8-10 investigated the hydroxyl radical reaction plays an insignificant role. Kinetics of Peroxone Oxidation The peroxone oxidation experiments were carried out at 25 °C to study the reaction kinetics in aqueous
Ind. Eng. Chem. Res., Vol. 36, No. 10, 1997 4111
solutions of pH varying from 8 to 10. The peroxone oxidation rate is enhanced in the presence of hydrogen peroxide at fixed initial concentrations of ozone and ammonia, and the depletion rate of ozone by the selfdecomposition is negligible. As illustrated in Figure 1 (case 3) and discussed earlier, the depletion rate of ozone is first order with respect to the ozone concentration. At fixed pH and temperature, the apparent rate constant, k′, can be plotted against the initial concentration of ammonia on a logarithmic scale for a series of experiments carried out at identical concentrations of hydrogen peroxide and ozone. The results reveal that the apparent rate constant increases with an ammonia concentration of the exponent of less than 0.2. This suggests that the peroxone oxidation is controlled mainly by the hydroxyl radical reaction of ammonia and that the depletion rate of ozone is influenced very little by the ammonia concentration. Under the above conditions, eq 15 can be simplified to
-d[O3]/dt ) k′[O3] ) a(1 + b)[H2O2][OH-][O3] (27) By plotting the apparent rate constant against the initial concentration of hydrogen peroxide on a logarithmic scale, a straight line can be drawn from the data obtained at fixed concentrations of ammonia and hydroxyl ion. The slopes for the individual lines are nearly unity at pH 8 and 9 and about 0.8 at pH 10. For the purpose of correlating the kinetic data to obtain an overall rate constant at a given pH, therefore, the reaction can be regarded as first order in ozone and hydrogen peroxide and independent of the ammonia concentration. On the basis of the above results, the depletion rate of ozone given in eq 27 can be rewritten in terms of the overall rate constant, k, as follows:
-d[O3]/dt ) k[H2O2][O3]
(28)
where k is equal to k′/[H2O2]0 for an experiment with excessive amounts of hydrogen peroxide. For the experiments carried out at identical pH and temperature, the individual rate constants are averaged to yield an average rate constant, as shown in Table 1. As stated earlier, the correlation coefficients are 0.98 or higher in obtaining apparent rate constants by regression analysis for the individual runs. The average deviations of the overall rate constants of the individual runs from the average values are 8.2, 10.6, and 19.4% respectively at pH 8, 9, and 10. Due to limitations of the stoppedflow apparatus, reliable data were not obtainable in solutions of higher than pH 10. For the peroxone oxidation, the theoretical rate expression for ammonia can be obtained from eq 16 by neglecting the rate of direct oxidation. The resultant equation can be combined with eqs 18 and 28 to yield
-d[NH3]/dt ) -c d[O3]/dt ) ck[H2O2][O3] (29) where c is the ratio of the ammonia depletion rate to the ozone consumption rate in the peroxone oxidation. The theoretical expression for c is
c ) k13[NH3]/{2k4[O3] + [k5 + k6(K2/K12)[OH-]][H2O2] + k13[NH3]} (30) The above equation suggests that the rate ratio, c, varies
depending upon relative magnitudes of the various concentrations, and equilibrium and rate constants. The equilibrium and rate constants at ambient temperatures (about 20 °C) have been reported as k4 ) 2.0 × 109 1/Ms (Buhler et al., 1984), k5 ) 2.7 × 10 7 1/Ms (Christensen et al., 1982), k6 ) 7.5 × 109 1/Ms (Christensen et al., 1982), k13 ) 1.0 × 108 1/Ms (Farhataziz and Ross, 1977), K2 ) 10-11.8 (Behar et al., 1970), and K12 ) 10-14, well-known in the literature. In the kinetic experiments for the peroxone oxidation of ammonia, the initial concentration of ozone was controlled at 1.5 × 10-4 M for most runs. The initial concentration of hydrogen peroxide varied from 9.80 × 10-4 to 1.95 × 10-2 M, and that of ammonia ranged from 6.80 × 10-3 to 4.34 × 10-2 M. At the lowest concentrations of hydrogen peroxide and ammonia, the initial c value is estimated to decrease from 0.52 to 0.48 as the pH increases from 8 to 10 on the basis of the above data. In the same pH range, the initial c value varies from 0.79 to 0.56 at the highest concentrations of hydrogen peroxide and ammonia. Since the ozone concentration decreases with time, the c value also increases during a reaction. The actual rates of consumption of both ozone and ammonia, however, decline rapidly with the reaction time because of rapid decrease in the ozone concentration. Discussion The kinetic results suggest that it is technically feasible to remove ammonia by the oxidation processes at room temperatures in alkaline solutions of pH ranging from about 8 to 11. The reaction rate is negligible in acidic solutions because of complete dissociation of ammonia to ammonium ion which is not reactive toward the oxidants. On the other hand, the dissolved ammonia exists mainly as free ammonia at about pH 11 or above, and there are little advantages in carrying out the treatment in solutions of higher pH. Although a free radical, NH2•, might be formed as an intermediate in the hydroxyl radical reaction of ammonia (Hoigne and Bader, 1978), only the final product, NO3-, was detected in this work. In a recent study of the peroxone oxidation of toluene in alkaline solutions, the ozone consumption rate was found to be nearly independent of the pollutant concentration (Kuo and Chen, 1996), as reported in this work for the peroxone oxidation of ammonia. The experimental and calculated results indicate that, at comparable conditions, the depletion rate of ozone is enhanced greatly by the presence of hydrogen peroxide. As shown in the kinetic model, ozone, hydrogen peroxide, hydroxyl ion, and others are competing with ammonia in consuming the hydroxyl radical formed in the solution. Thus, the oxidation rate of ammonia depends on the relative rates of the hydroxyl radical reactions of various species, as indicated in eqs 29 and 30. In applications of the peroxone oxidation process for wastewater treatment, the concentrations of ozone and hydrogen peroxide are often comparable though the ammonia concentration varies. When both concentrations of ammonia and hydrogen peroxide are very low compared with the ozone concentration, b approaches unity, and the rate ratio, c, or eq 30, is reduced to
c ) 0.5k13[NH3]/{k4[O3]}
(31)
The above equation implies that the ratio of the peroxone oxidation rate of ammonia to that of ozone is
4112 Ind. Eng. Chem. Res., Vol. 36, No. 10, 1997
proportional to the ratio of the rate of hydroxyl radical reaction of ammonia to that of ozone. Since k13/k4 ) 0.05 according to the literature information (Farhataziz and Ross, 1977; Buhler et al., 1984), the oxidation rate of ammonia may be very slow relative to the depletion rate of ozone. This may result in little improvement in the destruction rate of ammonia by applying the peroxone oxidation process. Though the consumption rate of ozone is accelerated greatly in comparison with the traditional ozonation process, it is consumed mainly as a scavenger in reacting with the hydroxyl radical formed. The benefit of addition of hydrogen peroxide for ammonia removal in the peroxone oxidation process may be realized for decontamination of wastewaters containing high concentrations of ammonia at pH 11 or less. If the concentration ratio, [NH3]/[O3], is larger than 40, the depletion rates of ozone and ammonia are comparable, as suggested by eqs 29 and 30. As the concentration ratio increases further, c approaches unity, indicating that the destruction rate of ammonia is nearly identical with the consumption rate of ozone. Under this circumstance, the peroxone oxidation process is effective and economical for the ammonia removal, and a high efficiency of ozone utilization can be expected. Effects of the hydroxyl ion concentration on the ozonation and peroxone oxidation rates can be examined by plotting the second-order rate constant (k14 or k) against the pH value, as shown in Figure 2. According to the kinetic model dictated by eq 19, the ozonation rate constant, k14, is independent of the hydroxyl ion concentration. Nonetheless, the average slope of 0.12 obtained in the semilogarithmic plot in the figure indicates that k14 increases slightly with the hydroxyl ion concentration. The slight dependence of the oxidation rate on the pH seems common in the ozonation of many organic and inorganic compounds (Langlais et al., 1991). The influence of pH on the depletion rates of ozone and ammonia is much more pronounced in the peroxone oxidation process. The figure shows that at 25 °C the overall rate constant increases with the hydroxyl ion concentration of the exponent of 0.71. A theoretical relationship between the overall rate constant and the hydroxyl ion concentration can be derived by comparing eq 29 with eqs 27 and 28.
k ) a(1 + b)[OH-]
(32)
In appearance, the above equation tends to suggest that the rate constant increases in direct proportion to the hydroxyl ion concentration (if a and b remain constants). However, b is a weak function of various concentrations, and reaction 6 in the kinetic model suggests that hydroxyl ions are also generated in the hydroxyl radical reaction of hydroperoxide ions. The production of hydroxyl ions in this reaction step may result in reduction of the consumption rate of the hydroxyl ions provided initially in the solution. Therefore, the increase in the overall rate constant with the hydroxyl ion concentration is less than linear, as observed by the experimental and calculated results. Conclusions A kinetic model has been proposed and rate equations derived for depletion of ozone and ammonia in aqueous solutions containing ozone or ozone-hydrogen peroxide
mixtures. The ozonation reaction is predominated by the direct oxidation of ammonia with ozone molecules. The slow depletion rate of ozone is first order with respect to both concentrations of ammonia and ozone, and the reaction rate increases slightly with pH and temperature. The peroxone oxidation of ammonia is controlled mainly by the formation of hydroxyl radical and its subsequent radical reactions with ammonia and other species. The depletion rate of ozone is first order with respect to both concentrations of ozone and hydrogen peroxide but is nearly independent of the ammonia concentration. As the pH increases from 8 to 10, the overall rate constant increases from 5860 to 133 000 1/Ms at 25 °C, indicating that the reaction rate increases in proportion to the hydroxyl ion concentration of the exponent of 0.71. The depletion rates of ozone and ammonia are comparable at high concentrations of ammonia in the solution. Under this circumstance, the peroxone oxidation process can be effective and economical (with a high efficiency of ozone utilization) for decontamination of wastewaters containing ammonia. Nomenclature a ) constant defined by eq 17 b ) function defined by eq 18 k ) overall rate constant, 1/Ms k′ ) apparent rate constant, 1/s kd ) rate constant for self-decomposition of ozone, 1/s kr ) rate constant for reaction r, 1/Ms Kb ) ionization constant Kr ) equilibrium constant for equilibrium reaction r n ) order with respect to ammonia concentration in ozonation r ) stoichiometric ratio of ozonation, mol of ozone/mol of ammonia t ) reaction time, s
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Received for review March 11, 1997 Revised manuscript received June 6, 1997 Accepted June 23, 1997X IE9702082
X Abstract published in Advance ACS Abstracts, August 15, 1997.
