Phsicochemical Treatment of Hazardous Wastes - Chapter 14 (end) pps

The only exception arose in the com-parison of Fenton’s reagent system with a combined Fenton’s reagent/ozonesystem, in which case there seems to be interference of ozone with Fenton’sre

or the reaction half times These investigations have been conducted underidentical experimental conditions so that treatment efficiency can be com-pared under the same conditions In addition, combinations of various AOPshave also been analyzed with respect to cost efficiency This comparison wascarried out by combining various oxidants such as oxygen, ozone, andhydrogen peroxide (H2O2) with catalysts such as ultraviolet (UV), Fe2+, andTiO2 In general, synergic effects exist when two oxidation systems werecombined This synergy was reflected in a marked increase in the free-radicalreaction pathway The effect was also observed to increase with the com-plexity of the oxidation systems used The only exception arose in the com-parison of Fenton’s reagent system with a combined Fenton’s reagent/ozonesystem, in which case there seems to be interference of ozone with Fenton’sreagent through oxidation of Fe2+, which reduces the amount of Fe2+ todecompose H2O2 into hydroxyl radicals.

14.2 Fundamental Theory

The destruction kinetics of organic pollutants may be inhibited in ters containing complex compounds Under such conditions, the combination

wastewa-534 Physicochemical Treatment of Hazardous Wastes

of UV, O3, and H2O2 systems could be a better process because hydroxylradicals are generated by several mechanisms As a result, UV/O3/H2O2 isless affected by color and turbidity in wastewater than UV/O3 or UV/H2O2system They are also applicable over a wider pH range when compared toUV/H2O2 systems Figure 14.1 shows a proposed mechanistic pathway thatexhibits a network of free-radical chain reactions involving HO2•/O2, HO3•/

O3•–, and OH•/O•– in the generation and consumption of OH• as proposed

by Hong et al (1996) This reaction network was a revised version based uponthe reaction pathways formulated by Staehelin and Hoigne (1985) and Peytonand Glaze (1988)

Kinetic rate constants of various reactions in the reaction pathways arelisted in Table 14.1 The model includes (1) dark reactions such as thedecomposition of O3 in water and the interaction of O3 and H2O2•, and (2)UV-assisted actions such as the photolysis of O3 and H2O2 In aqueousmedia, the free-radical chain reactions are initiated either by the k1 step

or by the k7 followed by the k6 step When H2O2 is added to ozonatedwater (i.e., peroxone), the k7 step that produces HO2 becomes inconse-quential, and a large amount of H2O2 is available through the addition ofthe peroxone system Next, the k6 step becomes the predominant super-oxide- and ozonide-producing pathway, an important chain-initiationreaction This is evident when the reaction rates of the k1 and k6 steps arecompared with a typical batch dose of H2O2 (e.g., 10 to 200 mg/L) Forapplied [O3] = 2 × 10–5 M (1 mg/L) and [H2O2]T = 1.5 × 10–3 M (50 mg/L)

at pH 7 (Hong et al., 1996):

k1[O3] [OH–] = 70(2 × 10–5)(10–7) = 1.4 × 10–10M–1 s–1 (14.1)

k6[O3] [HO2] = k6[O3] [H2O2]T ([H+]/K[H2O2] + 1 + K[H2O2]/[H+])–1

= (2.8 × 106)(2 × 10–5)(1.5 × 10–3)(2.5 × 10–5) = 2.1 × 10–6M–1 s–1 (14.2)Therefore, the enhancement of the peroxone system is due to faster chaininitiation In addition, when a large amount of H2O2 is added, the scavenging

of OH• by H2O2 (k5 step) may overtake the O3 (k4 step) For example, [O3] =

generation of hydroxyl radical will be presented in detail in Section 14.4.The reaction mechanism in Figure 14.1 demonstrates that pH is an importantparameter Due to the fact that some anions such as HO2 react faster thantheir conjugate acids, pH will influence the reaction kinetics and thus thesteady-state concentrations of various intermediates As a result, the treat-ment efficiency can be greatly impacted by pH The acid–base equilibria ofvarious reacting species must be considered, as shown in Table 14.1 (Hong

et al., 1996) The relative amounts of UV illumination, O3, H2O2, and

scav-engers (S i) such as HCO3 are critical in determining the level of active OH•radical in the system In UV-illuminated reactions, contaminant degradationrates are typically observed to be much faster than dark reactions (Zappi etal., 1993)

14.3 Process Description

From the late 1980s to the early 1990s, Ultrox International, a commercialmanufacturer, has demonstrated the efficacy of ultraviolet-light-enhancedoxidation at sites belonging to the Department of Defense (DOD) and theSuperfund sites (Zeff and Barich, 1992) Figure 14.2 illustrates a flow diagram

of the Ultrox UV/oxidation treatment system It shows that two differentoxidants are used in the process Ozone is generated from air and hydrogen

peroxide from the feed tank; therefore, the process can be operated in anycombinations, such as UV/O3, UV/H2O2, and UV/O3/H2O2.

Figure 14.3 presents an engineering drawing of the process Standardequipment designs is used in all of these installations Reactor size variesfrom 300 to 4800 gal, and reactors are fabricated from stainless steel Ozonegenerators range from 21 to 140 lb/day The ozone is dispersed throughporous, stainless-steel diffusers The number of diffusers required depends

on the organic compound being oxidized and the degree of removalrequired The UV lamps are enclosed within quartz tubes for easy replace-ment and are mounted vertically within the reactor (Zeff and Barich, 1992)

By utilizing a combined UV/O3/H2O2 process, the Ultrox oxidation tems proved to be efficient in removing volatile organic compounds(VOCs), benzene, toluene, xylene, hydrazines, phenols, chlorophenols,

OH • + A i = A′ i + OH

Source: From Hong, A et al., J Environ Eng 122(1), 58–62, 1966 With permission.

dioxanes, polychlorinated biphenyls (PCBs), and pesticides present inwastewaters and groundwaters

Figure 14.4 is the three-dimensional view of the process The figure trates the relative water head at different stages Because of flexible design,Ultrox UV/oxidation treatment systems have a number of advantages: (1)very few moving parts; (2) operation at low pressure; (3) minimum mainte-nance; (4) full-time or intermittent operation in either a continuous or batchtreatment mode; (5) use of efficient, low-temperature, and long-life UVlamps; and (6) use of a microprocessor to control and automate the treatmentprocess (Zeff and Barich, 1992)

illus-14.4 Degradation of Organic Pollutants

Phenol is degraded faster by ozone processes under basic pH than acidic

pH because the contribution of hydroxyl radicals increases with pH Table14.2 lists the degradation efficiency of phenol under various experimental

Ozone Diffuser (typical)

Treated Effluent Storage Tank

Effluent Sample Tap

variable conditions At neutral pH and at low H2O2 concentrations, H2O2improved ozonation slightly but showed an inhibitory effect at concentra-

tions higher than 6.2 mM The same effect was observed for the combination

of UV/O3/H2O2, and the limiting H2O2 concentration was found to be 0.07

mM Under UV illumination alone, the best conditions were found at pH 5

without any buffer The degradation rate increased considerably when H2O2was used; nevertheless, the initial H2O2 concentration exerted little influence

on the range used In photocatalysis, the degradation rate increased with thecatalyst concentration up to a value of 0.5 g/L From that point, the rate wasalmost constant In Fenton’s reaction, the limiting factor was the amount ofhydrogen peroxide The higher the amount of H2O2, the faster the degrada-tion will be Also, as the concentration of Fe(II) ion is increased, the degra-dation rate is improved

TABLE 14.2

Phenol Degradation in H2O2 Oxidation Systems as a Function of Time,

pH, and Initial Reactant Concentrations

Phenol Degradation (%)

O3

O3

O3

5.7–3 7.2 buffered 9.4 buffered

0 0 0

0 0 0

0 0 0

80 80 80

85.4 90.0 100

6.8 534 0.62 6.2 31 78 155 6.2 31

0 0 0 0 0 0 0 0 0

0 0 0 0 0 0 0 0 0

80 80 80 80 80 80 80 80 80

80.6 58.3 90.4 93.4 86.9 86.5 77.7 92.5 88.8

UV

UV

UV

4.4 – 4 6.8 buffered 11.5 buffered

0 0 0

0 0 0

0 0 0

30 30 30

24.2 14.0 5.0

0 0 0

0 0 0 0

30 30 30 30

24.2 87.1 90.6 89.8

5.2–3 6.9 buffered 9.4 buffered

0 0 0

0 0 0

0 0 0

80 80 80

80.9 92.6 91.9

Source: Esplugas, S et al., Water Res., 36, 1034–1042, 2002 With permission.

CH O

2

14.4.1.1 Comparison of Pseudo First-Order Kinetic Constant

The kinetic parameters chosen for comparison are rate constants and t1/2.Radiation influences and the effect of reactor design are usually identicalwhen these kinetic data are compared between the various AOPs tested Thevalues for pseudo first-order kinetics and half-lives for various processes are

given in Table 14.3 In most cases, the values of t3/4 are equal to two timesthose of t1/2; therefore, the reactions obey a first-order kinetics Figure 14.5.shows that Fenton’s reagent has the largest rate constant, e.g., approximately

40 times higher than UV alone, followed by UV/H2O2 and O3 in terms ofthe pseudo first-order kinetic constants Clearly, UV alone has the lowestkinetic rate constant of 0.528 hr–1

The economic evaluation of treatment processes is very important in ing an AOP from different available systems The overall costs are repre-sented by the sum of the capital, operating, and maintenance costs (Table

Comparison of the rate constants for selected AOPs All processes have been approximately

first-order with respect to substrate concentration (Data from Esplugas, S et al., Water Res., 36,

14.4) An estimation of costs has been made in this section; however, it should

be pointed out that costs could considerably decrease for photocatalytictreatments when solar light is used Figure 14.6 indicates that UV alone hasthe maximum cost because it has the lowest kinetic rate constant Fenton’sreagent has reasonably lower costs and the highest rate constant DifferentAOPs (ozone and its combinations, photocatalysis and UV/H2O2, photoca-talysis and Fenton’s reagent) have been compared in terms of the degrada-tion of phenol in aqueous solution In UV processes (UV, UV/H2O2, andphotocatalysis), the degradation rate produced by the UV/H2O2 process wasalmost five times higher than photocatalysis and UV alone Fenton’s reagentshowed the fastest degradation rate, 40 times higher than the UV processand photocatalysis and five times higher than ozonation Nevertheless, thedegradation rates and lower costs obtained with ozonation make it the mostappealing choice for phenol degradation

para-Hydroxybenzoic acid is a very common pollutant in a variety of

industrial wastewaters (olive oil and distillation industries) Because it is

highly toxic and refractory to anaerobic biological treatment, it was chosenfor the comparison as being representative of phenolic acids, and its deg-radation efficiency and kinetics are compared for several AOPs The firststage was to compare the 12 oxidation processes applied to the destruction

of p-hydroxy-benzoic acid Table 14.5 lists the semi-reaction times and the

conversions attained at 5 and 10 min of reaction time The kinetic eters of different oxidation processes are also shown in Table 14.5 Figure

param-14.7 demonstrates that Fenton’s reagent has the highest kinetic rates

fol-lowed by ozonation and then UV irradiation To study the improvement

by different AOPs, the various oxidation systems are divided into threegroups corresponding to the three basic oxidation processes from whichthey derive: UV irradiation, ozonation, and Fenton’s reagent

TABLE 14.5

Oxidation Rate Constants for Various AOPs

Oxidation Process k (min–1 ) t1/2 (min)

14.4.2.1 Oxidation Processes Using UV Radiation

The overall reaction process consists of three contributions: direct oxidation

by UV irradiation (photolysis), direct oxidation by ozone (ozonation), andoxidation by free radicals:

Direct oxidation of organic compounds by hydrogen peroxide in the entire

pH range is negligible At pH 5, ozone reacts with the compound directlyand not via free radicals (Staehelin and Hoigne, 1985) Figure 14.8 showsthat the order of efficiency in the oxidation systems under the UV irradiationcan be described as:

UV < UV/TiO2 < UV/O3 < UV/H2O2 < UV/H2O2/O3

The pseudo first-order kinetic constants (k t ) of p-hydroxybenzoic acid are

given in Table 14.6 UV/H2O2 has a higher rate constant than that of UV/

O3, because the former requires only 0.5 peroxide molecule and 0.5 photon,whereas the latter process requires 1.5 ozone molecules and 0.5 photon This

is shown in Equation (14.8) and Equation (14.9):

(14.8)(14.9)

Because H2O2 can be infinitely soluble in water, there is no limitation onmass transfer; however, O3 has a lower solubility in water and has masstransfer limitations Thus, the rate constant for UV/H2O2/O3 is the greatest,followed by that for UV/H2O2 and then UV/O3.

The pseudo first-order kinetic rate constants for processes based on theapplication of ozone are given in Table 14.7 Figure 14.9 presents the pseudo

first-order rate constants of p-hydroxybenzoic degradation The degradation

rates follow the increasing order:

O3 < O3/H2O2 < O3/H2O2/Fe2+ < O3/Fe2+ < UV/O3 < UV/O3/H2O2 < UV/

O3/H2O2/Fe2+

In regard to the degree of conversion of p-hydroxybenzoic acid at 5 and 10 min

of reaction time, the degradation efficient of AOPs using Fenton’s reagentfollows the increasing order:

TABLE 14.6

Oxidation Rate Constants for Various AOPs

Oxidation Process k t (min –1 )

Oxidation Rate Constants for Various AOPs

Oxidation Process k t (min –1 ) t1/2 (min)

Fe2+/H2O2/O3 < Fe2+/H2O2 < Fe2+/UV/H2O2 < Fe2+/H2O2/UV/O3

Table 14.8 shows the rate constants and reaction half times for oxidation of

p-hydroxybenzoic acid by these AOPs The results are plotted in Figure 14.10.

Figure 14.10 indicates that the addition of ozone has neither synergic noradditive effects (Fenton’s reagent + ozone) This may be because ozoneinterferes with the action of Fenton’s reagent, probably by oxidizing Fe2+ to

Fe3+ The combined effect is faster at the beginning (when Fe2+ concentration

is high), but is eventually overtaken by Fenton’s reagent The last two dation systems in Table 14.8 show the synergic effects of adding ozone tothe photo–Fenton’s reagent system due to the notable increase of degradation

oxi-rates The order of reactivity determined for the oxidation of

p-hydroxyben-zoic acid in aqueous solution is as follows:

UV < UV/TiO2 < O3 < O3/Fe2+ = O3/H2O2 < UV/O3 < UV/H2O2 =



 

Processes

14.4.3.1 Comparison of Various AOPs

The decomposition of all four selected CPs (4-CP, 2,4-DCP, 2,4,6-TCP, TeCP) was achieved by utilizing the following oxidants: UV radiation, Fen-ton’s reagent, and ozone The results are tabulated in Table 14.9 Figure 14.11

2,3,4,6-shows the degradation rate of chlorophenols by UV/H2O2 Figure 14.12

suggests that 4-CP was rapidly degraded by Fenton’s reagent, and ing rates were obtained for 2,4-DCP, 2,4,6-TCP, and 2,3,4,6-TeCP This sug-gests that the increase of chlorine atoms in a chlorophenol molecule decreasesthe susceptibility of aromatic rings to attack by the hydroxyl radicals gen-

decreas-FIGURE 14.10

p-Hydroxybenzoic oxidation rate constants for various Fenton’s reaction systems (Data from Beltan-Heredia, J et al Chemosphere, 42, 351–359, 2001.)

TABLE 14.8

Oxidation Rate Constants for Various AOPs

Oxidation Process t1/2 (min) k t (min –1 )

Comparison of Rate Constants for Various AOPs

Compound Ultraviolet Fenton O 3 at pH = 2 O 3 at pH = 9

erated by UV radiation This can be explained by the electronically excitedstates of polychlorinated phenols generated during photochemical treat-ments (Burrows et al., 1998) In these excited states, the CP moleculesundergo intramolecular transformations and stabilize with different electrondistributions, followed by decomposition to radical or molecular products.

A higher degree of chlorine substitution withdraws electrons from the matic rings As a result, electrophilic addition of hydroxyl radical to the ring

aro-is not favored The formation of an excited state or stabilization of theintermediate state becomes more difficult to oxidize than non-chlorine sub-stituted phenols

Figure 14.13 shows that the degradation by ozone alone is much faster at

pH 9 than at pH 2 In addition, the greater the number of chlorine uents contained in a chlorophenol molecule, the higher the degradation rate

substit-is The figure indicates that the oxidation rates are not the same as in thephotodegradation process In UV/O3, hydroxyl radicals usually attack thearomatic ring at the sites that are not occupied by chlorine atoms; therefore,

FIGURE 14.11

Benitez, F.J et al., Chemosphere, 41, 1271–1277, 2000.)

FIGURE 14.12

Decomposition rates for selected chlorophenols in the Fenton’s reaction system (Data from

Benitez, F.J et al., Chemosphere, 41, 1271–1277, 2000.)

0 100 200 300 400 500 600

CPs

0 200 400 600 800 1000 1200 1400 1600 1800 2000

CPs

hydroxylation is the first elementary step and precedes dissociation of rine atoms (Tang and Huang, 1996) So, the increase in number of chlorineatoms in the aromatic ring usually decreases the reactivity toward hydroxylradicals However, at pH 2 and 9, hydroxyl radical is not the dominantspecies in ozonation alone Therefore, oxidation by molecular ozone contrib-

chlo-utes mostly to the CP degradation The trend in kp values is reversed and kpincreases with increasing substituent chlorine atoms The rate of decompo-sition increases in the following order as shown in Figure 14.13:

2,3,4,6-TeCP < 2,4,6-TCP < 2-4-DCP < 4-CPThe pathways proposed by Chen and Ni (2001) to describe the completeozonation of 2–chlorophenol were hydroxylation, dechlorination, and ringcleavage According to this mechanism, the presence of more atoms of chlo-rine enhances the dechlorination step; therefore, the degradation is faster.Degradation efficiency of CPs can be compared for combinations of UVradiation + H2O2, UV radiation + Fenton’s reagent, and UV radiation + ozone

as described in the following sections

Photodegradation experiments with the four selected CPs in the presence

of hydrogen peroxide were conducted at 25°C and pH 2 by Benitez et al.(2000) The degradation rates achieved were the same as those obtainedfor the single photochemical process The degradation rates decreasedwhen the number of chlorine substituents increased, as shown in Table14.10 Figure 14.14 shows the combined process constant (k t) and radical

rate constant (k r) It can be seen that these rate constants have moderatelyhigher values than those obtained in the single photodecomposition pro-cess due to hydroxyl radical reactions

FIGURE 14.13

Effect of pH on chlorophenol oxidation rates at various pH values (Data from Benitez, F.J et

al., Chemosphere, 41, 1271–1277, 2000.)

0 50 100 150 200 250 300 350 400 450

CPs

pH 2 pH9

14.4.3.3 Photo–Fenton’s Reagent System

Decomposition experiments for these CPs listed in Table 14.11 were carriedout by the simultaneous action of UV radiation and Fenton’s reagent (Benitez

et al., 2000) Table 14.11 shows the first-order rate constants and half-lives.During the photo–Fenton’s reagent reaction, the single photodecomposition

rate constant, k t, decreased as the number of chlorine substituents increased

In addition, combined rate constants, k t, are much greater than the radical

reaction constants, k r Therefore, this confirms the additional contribution ofthe radical reaction due to generation of the hydroxyl radicals by Fenton’s

TABLE 14.10

Oxidation Rate Constants for Various AOPs

Compound k t× 10 3 min –1 k r× 10 3 min –1

Benitez, F.J et al., Chemosphere, 41, 1271–1277, 2000.)

TABLE 14.11

Oxidation Rate Constants for Various AOPs

Compound k t × 103 min –1 k r × 103 min –1 t1/2 (min)

4-CP 2,4-DCP 2,4,6-TCP 2,3,4,6-TeCP

Chloro compounds

reagent and by UV/H2O2 Figure 14.15 shows the radical decomposition rateconstants for the photo–Fenton’s reagent reaction Figure 14.16 plots the com-bined rate constants for the photo–Fenton’s reagent system It show that the

overall rate constants, k t , increased much more than the radical constants, k r

Finally, degradation experiments with these four CPs by the combined cess UV/O3 were conducted at 25°C and pH 2 by Benitez et al (2000) Table14.12 shows the rate constants, k t, and determined through kinetic studies

pro-Figure 14.17 shows that the rate constants decreased significantly whenchlorine substituents were increased from one to two However, the rateconstants remained relatively the same for 2-, 3-, and 4-chlorine-substitutedphenols Comparing these results to those of the single photodecomposition

or pH 2 ozonation processes, this combination accelerates the decomposition

rate, with an extremely high rate constant (k t) for 4-CP and lower rates forthe other CPs The remaining CPs (2-4-DCP, 2,4,6-TCP, 2,3,4,6-TeCP)