Phsicochemical Treatment of Hazardous Wastes - Chapter 6 pps

According to the slopes of the straightline, the rate constants between hydroxyl radicals and 2,4-DCP and 2,4,6-TCP can be determined as 7.22 × 109 M–1 s–1 and 6.27 × 109 M–1 s–1, respec

of the organic acids (Fenton, 1899, 1900) In addition, Cross et al (1900)further confirmed that ferrous salts significantly enhance the kinetics ofhydrogen peroxide decomposition Goldhammer (1927) investigated theeffect of Fenton’s reagent on phenols and found that for each equivalent of

Fe2+ three equivalents of H2O2 were decomposed They also noted that inconcentrated hydrogen peroxide solutions each mole of Fe2+ decomposed 24equivalents of hydrogen peroxide

Haber and Weiss (1934) were the first to propose that free radicals existed

as intermediates during the chemical reactions in solution The next year,Haber and Weiss further investigated the Fenton chemistry and concludedthat Fenton’s reaction can be expressed as a series of chain reactions withreaction pathways dependent on the concentration of the reactants Thestudy disproved the original theory of Fenton’s reaction, which suggestedthat the interaction between an intermediate, six-valent, iron–oxygen com-plex and hydrogen peroxide was the most significant reaction step In 1934,Haber and Weiss proposed that breaking rate of chain length wasincreased at lower pH so the propagation cycle was extended before termi-nation The concentration of free hydroxyl radicals was determined to bedirectly proportional to the concentration of hydrogen peroxide

Baxendale and Wilson (1957) reported that in an oxygen-free environmentFenton’s reagent initiates very rapid polymerization of methyl acrylate,methacrylic acid, methyl methacrylate, acrylonitrile, and styrene, and the

166 Physicochemical Treatment of Hazardous Wastes

reaction is a function of the concentration of hydroxyl radicals In the ence of oxygen, no polymerization occurs Barb et al (1951) conducted anextensive investigation of Fenton’s reagent chemistry When [H2O2]/[Fe2+]ratios are low, the reaction rate is second order and stoichiometry is2[Fe2+]≅[H2O2]; however, in the presence of polymerizable vinyl compoundthe reaction remains second order but the stoichiometry changes to[Fe2+]≅[H2O2] Thus, they concluded that polymerization of vinyl compoundsoccurs and results in a polymer with terminal hydroxyl groups An inhibitioneffect of hydroxyl radicals due to the higher concentration of hydrogenperoxide was also suggested To explain this mechanism, it was proposedthat hydroxyl radicals react with hydrogen peroxide to form hydrogen diox-ide This process decreases the hydroxyl radicals generated by the reactionbetween ferrous iron and hydrogen peroxide In addition, Barb et al (1951)suggested that hydrogen dioxide is not a strong oxidizing agent capable ofbreaking the bonds of vinyl compound or oxidizing other organics Merz and Waters (1949) showed that oxidation of organic compounds byFenton’s reagent could proceed by chain as well as non-chain mechanisms,which was later confirmed by Ingles (1972) Kremer (1962) studied the effect

pres-of ferric ions on hydrogen peroxide decomposition for Fenton’s reagent Itwas confirmed that once ferric ions are produced the ferric–ferric system iscatalytic in nature, which accounts for relatively constant concentration offerrous ion in solutions

In the late 1970s, two major theories were considered: the free radicalmechanism by Walling and Cleary (1977) and complex formation by Kremerand Stein (1977) Walling proposed that Fenton’s oxidation predominantlytakes place by the free-radical mechanism On the other hand, Kremer pro-posed that complexation between the iron and the organic molecules has asignificant role and thus concluded that both mechanisms occur simulta-neously In the late 1980s a simultaneous effort was made to apply Fenton’sreagent to the field of environmental science Various contaminants werestudied in the laboratory to determine the optimum conditions Practicalapplications of Fenton’s reagent to treat contaminants have also been exam-ined by pilot-plant and continuous treatment systems in textile wastewater,etc.; for example, Bigda (1996) applied Fenton’s reaction to the design of areactor for treatment of organic contaminants

6.2 Kinetic Models

Although Fenton (1894) studied the violet color in caustic alkali duringoxidation of tartaric and racemic acids by ferrous salt and hydrogenperoxide, no reaction kinetic model was offered Fenton reported that thecolor disappeared when acid was added Also, it has been observed thatfresh external air is more active than room air Fenton performed differentexperiments using various amounts of ferrous and hydrogen peroxides and

Fenton’s Reagent 167

proposed that iron catalyzed this reaction For example, a small amount ofiron is sufficient to determine oxidation of an unlimited amount of tartaricacid In tartaric acid, two atoms of hydrogen are removed from a molecule

of acid, resulting in the production of dihydroxymaleic acid Among mon oxidants such as chlorine, potassium permanganate, atmospheric oxy-gen, and electrolysis, the most effective oxidizing agent is hydrogenperoxide Fenton’s work was extended to alcohols (Fenton, 1899) and otherorganic acids (1900) Attempts to identify the intermediates and products ofseveral organic acids and alcohols were made without success

Goldschmidt and Pauncz (1933) suggested that Fenton’s reaction is a chainreaction involving the same reactive intermediates occurring during catalyticdecomposition of H2O2 rather than via formation of peroxides of iron:

It was also shown that the ratio of oxidized alcohol to oxidized Fe2+ could

be greater then one Baxendale and Wilson (1957) showed that hydroxylradical initiating the chain polymerization of olefins by hydrogen peroxidewas the same process as the rapid oxidation of glycolic acid Merz and Waters(1947) confirmed that simple water-soluble alcohols are oxidized rapidly byFenton’s reagent The primary alcohols are oxidized to aldehydes, which arefurther oxidized at comparable rates by exactly the same mechanism Merzand Waters proposed a mechanism of chain oxidation of alcohols and alde-hydes by sodium persulfate, hydrogen peroxide, and an excess of ferroussalt as follows:

4 Chain ending at high substrate (alcohol) concentration:

2R–CHOH• = R–CHO + R–CH2OH (disproportionation) (6.6)

168 Physicochemical Treatment of Hazardous Wastes

In 1949, Merz and Waters determined the values for the ratio of rateconstants k2/k3 that indicated which particular radical reduced hydrogenperoxide Based on the reaction pathways, they classified the reacting com-pounds into two groups The first group of substrates reacts by chain process.Only a small amount of reducing agent is required The second group iscomprised of substrates that react by non-chain processes — in this case, theoxidation is caused by the hydroxyl radical, and considerable loss ofhydroxyl radical occurs For the first group, the reaction rate can be expressed

by Equation (6.7):

d[H2O2]/d[RH] = 1 + k2[Fe2+]/k3[RH] (6.7)For non-chain reactions, the kinetic rates are described by Equation (6.8):

d[H2O2]/d[RH] = 2 + k2[Fe2+]/k3[RH] (6.8)The values for the ratio of rate constants k2/k3 can be determined from theintercept of their graphs The results will suggest which particular radicalreduced hydrogen peroxide

Barb et al (1951) gave a redox formulation that involves the followingreaction sequence:

HO•2+Fe2+k=5Fe3++OH2–

OH•+Fe2 +k=6Fe3 ++OH–

Fenton’s Reagent 169

where k1 and k2 showed inverse [H+] dependence

The following scheme was presented by Kremer and Stein (1959), and further

elaborated by Kremer (1963):

(6.15)

(6.16)FeO3+ + H2O2 = Fe3+ + H2O+ O2 (6.17)Let C1 = [H+] and C2 = [FeO3+], k aand k d showed inverse [H+] dependence

and k b >> k a >> k c, C1 could be taken as a low concentration intermediate to

a good approximation

[C1] = K[H2O2][Fe3+], K = k a/k b (6.18)[Fe3+]t= [C2] + [Fe2+] (6.19)–d[H2O2]/dt = k c K[Fe3+]t[H2O2] + (k d – k c K)[C2][H2O2] (6.20)

d[O2]/dt = k d[C2][H2O2] (6.21)

d[C2]/dt = k c K[Fe3+]t[H2O2] – (k d + k c K)[C2][H2O2] (6.22)

[C2] rises continually during the reaction, approaching a saturation value

of k c K[Fe3+]t /(k c K + k d), and –d[H2O2]/dt is always greater than twice d[O2]/

dt At the end of the reaction, some hydrogen peroxide will be stored as

C2, and less than 0.5 mol of O2 will be liberated per mole of H2O2

decom-posed

Walling and Kato (1971) modified the reaction mechanism proposed by Merz

and Waters as follows:

Fe2+ + H2O2 = Fe3+ + OH– + OH•, k1 = 76 (6.23)

Fe3++H O2 2⇔k k FeOOH2++H+

b a

FeOOH2+k=cHO–+FeO3+

where k is in L/mol/s, taken from the literature The reaction conditions

were chosen to minimize the competing processes as follows:

HO• + H2O2 = H2O + HO2, k = (1.2–4.5) × 107 (6.31)

2HO• = H2O2, k = 5.3 × 109 (6.32)Thus, the stoichiometry is:

where R = ∆[Fe2+]/2∆[H2O2], a = k2/Σk3, r = [Fe2+]/2[RH], and b = (k 3j + 2k 3k)/2Σk3 This mechanism is referred to as the free-radical mechanism.

In 1972, Ingles reported his studies of Fenton’s reagent using redox titration

He found evidence in support of Kremer’s complex mechanism theory andconcluded that, when suitable complexes are formed, substrates are notoxidized by free radical; rather, electron transfer processes might be

involved Fenton’s reaction scheme was modified by Ingles for the case whensubstrate is present in large amounts in the form of substrate/iron-peroxidecomplexes Ingles suggested that electron transfer occurs within this com-plex

(6.34)All substrates were considered to compete as ligands in iron complexesand to modify the reaction characteristics of each other and of the complex.Reaction 6.34 yields hydroxyl radicals, so the free-radical mechanism pro-posed by Walling appeared to be possible; however, Equation (6.35) to Equa-tion (6.38) involve electron transfer and do not lead to formation of hydroxylradicals Equation (6.37) and Equation (6.38) involve ionic mechanisms:

by Tang and Huang (1996a) The equation used to calculate the rate constants

rad-[S] = concentration of the substrate at any time.

[S0] = initial concentration of the substrate

[R] = concentration of the reference compound at any time

[R0] = initial concentration of the reference compound (2-chlorophenol)

In their work, the reference compound is 2-chlorophenol, with a rate constant

of 8.2 × 109 M–1 s–1 Either 2,4-DCP or 2,4,6-TCP was mixed with nol in a reactor, separately Then, H2O2 was mixed with the organic com-pound and the pH was adjusted to 3.5 The organic concentrations weremeasured by gas chromatography (GC) before and after Fe2+ was added.The results are shown in Figure 6.1 According to the slopes of the straightline, the rate constants between hydroxyl radicals and 2,4-DCP and 2,4,6-TCP can be determined as 7.22 × 109 M–1 s–1 and 6.27 × 109 M–1 s–1, respectively.The hydroxylation rate constants for 2,4-DCP and 2,4,6-TCP are clearlysmaller than that for 2-chlorophenol; therefore, increasing chlorine content

2-chlorophe-on the aromatic ring decreases the reactivity of the chlorinated phenolstoward hydroxyl radical attack

6.2.6.2.1 Pseudo First-Order Kinetic Model

When an excess of H2O2 and Fe2+ is added at constant concentrations to thesystem, a steady-state concentration of hydroxyl radical can be assumed.The concentration of both H2O2 and Fe2+ can be considered as constant;therefore, the pseudo first-order kinetic can be developed as follows:

Chlorinated phenols + •OH = intermediates (chlorinated aliphatic compounds) (6.40)

where: k1 is the pseudo first-order rate constant of oxidation

where k2 is the pseudo first-order rate constant of dechlorination The radation kinetics can be modeled as the following:

(Cl)/(CP)0 = 1 + [k1exp(–k2t) – k2exp(–k1t)]/(k2 – k1) (6.46) where (CP)0 is the initial concentration Figure 6.2 shows both the experi-mental data and the concentration profile predicted by the kinetic model forthe oxidation and dechlorination of 2,4,6-TCP

FIGURE 6.1

0 0.5 1 1.5 2 2.5

It is important to note that both H2O2 and Fe2+ have to be overdosed tomaintain a steady-state concentration of hydroxyl radical and to obtain asatisfactory approximation of the mathematical model with the experimen-tal data When H2O2 and Fe2+ concentrations are 5 × 10–3 M and 2 × 10–4 M,

respectively, the relative rate constants of 2-chlorophenol (2-CP) and TCP with respect to 2,4-DCP can be calculated The oxidation and dechlo-

2,4,6-rination constants of 2,4-DCP were found to be 0.995 1/min (k1) and 0.092

1/min (k2), as reported in a previous study (Tang and Huang, 1996) Forcomparison, Table 6.1 summarizes all the kinetic constants as determined

in this study and in the related literature

FIGURE 6.2

Kinetic modeling of 2,4,6-trichlorophenol oxidation and chloride ion dissociation (in the

TABLE 6.1

Kinetic Rate Constants of Chlorinated Phenols by Fenton’s Reagent

Measured oxidation constants

Cl-To evaluate the effect of the number of chlorines on the degradation rateconstants of different chlorophenols, Table 6.2 shows the rate constants of

elementary, oxidation, and dechlorination for the ratios of k2-CP/k2,4-DCP and

k2,4,6-TCP/k2,4-DCP The relative rate constants are plotted against the number ofsites unoccupied by chlorine atoms on the chlorinated phenols in Figure6.3.A linear correlation between the rate constants and the number of sitesavailable is found with a standard deviation of 0.132 Clearly, the morechlorine atoms the aromatic rings contain, the fewer sites are available forhydroxyl radical attack; however, the correlation should not be used for

TABLE 6.2

Relative Ratios of Kinetic Constants Using 2,4-DCP as the Reference Compound

(k/k2,4-DCP)

FIGURE 6.3

The correlation between oxidation constants and sites available on aromatic ring for hydroxyl

y=0.755X - 1.328

00.511.52

Number of Sites Available

relative rate constants measured for different chlorophenols correlation of the rate constants of chlorophenol

predicting the rate constants of tetra-chlorophenol and penta-chlorophenol

due to steric hindrance Figure 6.3 indicates that the oxidation rate decreaseswith the increasing degree of chlorine content in the following order (interms of both elementary rate constants and the observed pseudo first-orderrate constants): 2-CP > 2,4-DCP > 2,4,6-TCP

6.2.6.2.2 Dechlorination Kinetic Model Using Transition State Theory

When both H2O2 and Fe2+ are not overdosed, the concentrations of both

H2O2 and Fe2+ will change The pseudo first-order kinetic model developedabove does not apply In order to quantitatively model the effect of H2O2and Fe2+ on the dechlorination kinetics by Fenton’s reagent, the dechlorina-tion kinetic model is developed as follows First, hydroxyl radicals are gen-erated by H2O2 decomposition by Fe2+:

be released from the chlorinated aliphatic intermediates instead of the matic ring When chlorinated phenols are present in the system, the follow-ing reaction mechanisms can be assumed:

b f

(CP OH)− *→k mchlorinated aliphatic intermediatesChlorinated aliphatic compounds OH+• →k p CO2+Cl–+H O2

attack on chlorinated phenols; k m is the rate constant for the formation of

intermediates; and k p is the rate constant for the formation of products.Because of the high reactivity of hydroxyl radicals, activated complex, andchlorinated intermediates, their concentrations are extremely low at thesteady state; therefore, a pseudo first-order steady state can be assumed forthe kinetic modeling As a result, the steady-state concentration of the acti-vated complex can be obtained by setting the change of its concentration tozero:

(6.54)

where (C*) is the concentration of the activated complex

Thus, the steady-state concentration of the activated complex should be:

(6.55)

For chlorinated aliphatic intermediates, the steady-state concentration can

be derived by the same principle:

(6.56)Then, the concentration can be expressed as follows:

(6.57)

Similarly, the change of hydroxyl radical concentration should also equalzero:

(6.58)Substituting Equation (6.55) into Equation (6.58), we obtain:

Simplifying this equation, we get:

(6.60)

For simplicity, we assume that the following expression containing rate

constants is also a constant k:

mech-rophenols It is not certain, however, that it can be applied to tetra- and

penta-chlorophenols due to steric hindrance; therefore, when the above generalequation is applied to chlorophenols, the equation becomes:

2(H O ) (CP)

0

The k can be obtained using the initial rate method by overdosing H2O2 Forexample, at the overdosed H2O2 concentration of 5 × 10–3 M and constant

pH of 3.5, the k can be obtained by plotting the 1/r0 vs 1/Fe2+ during thedechlorination of 2,3,4-trichlorophenol

Figure 6.4 shows that the slope of the line is 7.21 (s) As a result, the k can

is about the same as when H2O2 was overdosed

Using the developed model, the k values for 2-CP, 3-CP, and 4-CP are 1.12

× 107, 1.004 × 109, and 1.005 × 108(1/s), respectively; therefore, the nation constants for monochlorophenols follow a decreasing order: 3-CP >4-CP > 2-CP Because chloride ion can be released only after the rupture ofthe aromatic ring, the faster the hydroxylation of the parent compounds, thefaster the dechlorination process should be Therefore, the above order can

dechlori-be understood in terms of the effect of the substituents on the reactivity of

their parent compounds It is known that both OH and Cl are ortho and para

directors Under the influence of these directors, the following preference ofhydroxyl radical attack is expected:

FIGURE 6.4

y=7.21x + 1.94 * 10 4

0 20000 40000 60000 80000 100000 120000 140000 160000 180000 200000

t i

In Relation (6.67), the solid arrow presents a stronger directory effect bythe hydroxyl group than that by the chlorine group, which is presented by

the dash arrow It can be seen that 3-CP has three ortho and para positions

enhanced by OH and Cl directors For 4-CP and 2-CP, however, no position

is enhanced by both OH and Cl directors Because of these directors, mediates with higher degrees of oxidation are expected to be produced inoxidation of 3-CP compared to the oxidation of 4-CP and 2-CP Therefore,the dechlorination rate constants of 4-CP and 2-CP will be smaller than that

inter-of 3-CP On the other hand, 2-CP will have some steric hindrance effect due

to the OH and Cl groups on the aromatic ring being located closer than isthe case for 4-DCP As a result, 2-CP will be more difficult to oxidize than

4-CP In other words, because the ortho position of chlorine is closer to the hydroxyl group than the meta and para positions, it will be subject to more

steric strain than other congeners and have a greater change of free energy

after dechlorination; therefore, hydroxylation at the ortho position will rience more steric strain than meta or para monochlorophenols

H2O2/(CP) ratio during the oxidation of 2-MCP, 2,4-DCP, and 2,4,6-TCP at

a constant H2O2/Fe2+ ratio of 2.5 and optimal pH of 3.5 When the ratio of

H2O2/DCP increases, the dechlorination rate constants increase more, the difference between dechlorination rate constants becomes more

Further-FIGURE 6.5

y=54.1x + 1380

0 50000 100000 150000 200000

and more apparent when the H2O2/DCP ratio increases from 1 to 20 For aconstant H2O2/CP, the H2O2/Fe2+ ratio is constant Therefore, the steady-state concentration of hydroxyl radicals is constant However, as the ratio of

H2O2/DCP increases, more hydroxyl radicals are available to react with CP.Because chloride ion is released through numerous hydroxylation andhydrogen abstraction steps, more hydroxyl radicals are available for a DCPmolecule, and the faster the dechlorination kinetics will be The effect of the

H2O2/CP ratio on the dechlorination constants can be clearly seen in Figure

6.6 The dechlorination rate follows this decreasing order: kTCP > kDCP > kMCP

At the optimal H2O2/Fe2+ ratio, the amount of H2O2 available for 1 mol ofchlorinated phenols affects dechlorination rate constants significantly, asshown in Figure 6.9 When the ratio of H2O2/CP is 20, the dechlorinationconstants for 2-CP, 2,4-DCP, and 2,4,6-TCP are 0.05, 0.16, and 0.33 (1/s),respectively The relative dechlorination constants of these chemicals are1:3:6 At a constant H2O2/CP ratio, the maximal capacity of Fenton’s reagent

is constant, because the number of available sites not occupied by chlorineatoms is the major factor responsible for dechlorination If we assume thateach site has an equal probability of being attacked by hydroxyl radicals,the efficiency of dechlorination by 1 mol of hydroxyl radicals should be 1/

4, 2/3, and 3/2 for 2-CP, 2,4-DCP, and 2,4,6-TCP, respectively This gives1:2.6:6 as a relative dechlorination rate constant The theoretical predictionagrees fairly well with the experimental result of 1:3:6 This is the reasonwhy no linear correlation could be found between dechlorination constantsand the nonchlorinated sites available All the sites unoccupied by chlorine

FIGURE 6.6

0 0.05 0.1 0.15 0.2 0.25 0.3 0.35

atoms on the aromatic ring of the chlorinated phenols have the same tivity toward hydroxyl radicals when chlorine atoms occupy the 2, 4, and 6positions Table 6.3 shows that the effect of the chlorine position on dechlo-rination rate constants decreases according to the following order: 2,5-DCP

reac-> 3,5-DCP reac-> 2,3-DCP reac-> 2,6-DCP reac-> 2,4-DCP

The ortho and para director nature of Cl seems to play a less important role

when the number of chlorine atoms increases from one to three; tion has the following order in terms of decreasing rate constants: 2,4,6-TCP

dechlorina-> 2,4,5-TCP dechlorina-> 2,3,4-TCP This order suggests that the steric hindrancebecomes a determining factor in preventing hydroxyl radical attack on theunoccupied sites of aromatic rings For example, 2,4,6-TCP has much largerspace than 2,4,5-TCP due to separation of the chlorine atoms on the aromaticring 2,4,5-TCP, in turn, has a larger space than does 2,3,4-TCP in whichchlorine atoms can locate closely to one another Therefore, 2,4,6-TCP isaffected the least by steric hindrance; 2,4,5-TCP is subjected to an averagesteric hindrance effect; and 2,3,4-TCP is subjected to the greatest steric hin-drance The dechlorination rate constants seem to follow the same order

suggested by steric hindrance The nature of the ortho and para directors of

Cl group seems to have diminished to a certain degree; in other words, ifthe directory effects of Cl group are predominant, then the order of decreas-ing dechlorination rate constants should be somewhat reversed

From Equation (6.64), the following conditions have to be valid so that theinitial dechlorination rate is independent of the organic concentration:

k(CP) >> [k t2(Fe2+) + k t3(H2O2)] (6.68)Under this condition, the numerical value of this expression is 5.1 × 10–4

(M–1 s–1), which is one magnitude larger than the experimental average value

of 5.1 × 10–5 (M–1 s–1), as shown in Figure 6.7 Nevertheless, this implies thatthe limiting step in the oxidation of chlorinated phenols is the generation ofhydroxyl radicals through Fenton’s reagent

2,3-Dichlorophenol 8.61 2,5-Dichlorophenol 11.90

r = k0 i(H O ) (Fe2 2 0 2+)0

6.2.6.2.3 Oxidation Model of Unsaturated Aliphatic Compounds

The same transition complex approach and steady-state assumptions wereused to develop the kinetic model of unsaturated chlorinated aliphatic com-pounds such as trichloroethylene (TCE) The model reflects the effects of

H2O2, Fe2+, and organic compounds on the oxidation kinetics as follows:

by keeping two variables constant and varying one variable among theconcentrations of H2O2, Fe2+, and substrate When the logarithm of both sides

of Equation (6.70) is taken, the following linear relationship can be obtained:

Using the experimentally determined constants such as ki, k+2, and k+3, the

kinetic models for dichloroethylene (DCE), TCE, and tetra-CE can be

dichloroethane oxidation does not give a straight line between log(r0) and

log(X), because the oxidation pathway of dichloroethane is different from

that in oxidizing chlorinated ethylenes For example, dichloroethane is one

log(r ) = log(k ) + n * log k log

of the saturated chlorinated aliphatic compounds, and the first oxidativestep is hydrogen abstraction instead of hydroxylation (Walling, 1975) Theunsaturated bond formed after hydrogen abstraction is then attacked by theaddition of hydroxyl radicals As a result, hydrogen abstraction must precedehydroxylation in the oxidation of dichloroethane.

In order to obtain the optimal ratio of H2O2 to Fe2+, Equation 6.76 can bedifferentiated with respect to H2O2, assuming that Fe2+ is the optimal con-centration of (Fe2+)opt We set:

(6.78)

The optimal concentration of H2O2 can be expressed as follows by solvingthe above equation

(6.79)

The optimal concentration of Fe2+ can be derived from the same

mathemat-ical approach We set the derivative of r s with respect to Fe2+ concentrationequal to zero:

(6.80)

Then, the optimal concentration of Fe2+ is:

(6.81)

When Equation (6.79) is divided by Equation (6.81), the optimal ratio of

H2O2/Fe2+ can be obtained:

s

r =

s

r =

(Fe )(H O )

2 2 2+

2 3

2+

2 2

= (k ) (k ) *

2 2 2

2 3

k k

Substituting the numerical values into the above equation, we obtain theoptimal ratio of H2O2 to Fe2+ as the following:

(6.84)

Experimental data indicate that the optimal ratio of H2O2 to Fe2+ is about

5 to 11 To investigate the effect of a low H2O2/Fe2+ ratio, the Fe2+ tration was overdosed at a constant value of 10–3 M At a high H2O2/Fe2+ratio, where H2O2 was overdosed at 10–2 M, the extrapolated maximum initial rate of 1.8 mM/min also occurs at an H2O2/Fe2+ ratio of 11 This valuereasonably agrees with the theoretical value of 11 as the optimal ratio of

concen-H2O2/Fe2+ At a constant H2O2/Fe2+, the H2O2 concentration required for percentage release of chloride ion in dichloroethylene is shown in Figure6.8.Tang and Huang (1997) concluded that the amount of H2O2 required for

iso-a specific percentiso-age removiso-al of the orgiso-anic compounds depends upon theinitial organic concentration to be oxidized This is also true for the totalpercentage of chloride ion released at different initial organic concentrations.The typical percentage removal of organic compounds and percentagerelease of chloride ion have been studied at 100, 70, 50, 40, 30, 20, 10, and1% The amount of H2O2 required to achieve a certain percentage removal

follows the order of TCE < tetra-CE < DCE << DCEA (dichloroethane) at a

Fe2+ concentration of 10–3 M However, the amount of chloride ion detected

at an H2O2 concentration of 10–2 M follows the order of DCEA << DCE < TCE < tetra-CE It is much more difficult to remove chloride atoms from

saturated aliphatic compounds such as DCEA than from unsaturated phatic compounds

ali-6.3 Oxidation of Organic Compounds

Trihalomethanes (THMs) are priority pollutants listed by the U.S mental Protection Agency (EPA) They are recalcitrant in nature, thus theirdestruction is difficult The most commonly encountered THMs in drinkingwater threatening human health are chloroform, bromodichloromethane,dibromochloromethane, and bromoform Tang and Tassos (1997) studied theoxidation kinetics and mechanisms of these four THMs

Environ-The effect of the ratio of H2O2 to Fe2+ on oxidation kinetics, the oxidationkinetics of THM mixtures, and the effect of the number of chlorine atoms in

a THM on its oxidation were all investigated Bromoform is the easiest tooxidize of the four THMs Bromoform concentrations used in the study ofFenton’s reagent ratio and oxidation kinetics were 49.2, 98.3, and 295 µg/L

As the ratio of H2O2 to Fe2+ increases, the removal efficiency increases with

(H O )(Fe )

3*10

2 2 2+

2 3

8 7

= k

k = =

the initial concentration of bromoform This may indicate that the hydroxylradical has a preference toward organic compounds, resulting in proportion-ately less scavenging effect by H2O2:Fe2+ and MeOH For a higher H2O2:Fe2+ratio of 10:1, the amount of bromoform removal appeared to show depen-dence on initial bromoform concentration As the initial organic substrateconcentration increases, less scavenging of OH• occurs At an H2O2:Fe2+ ratio

of 100:1, bromoform removal was only 25%, while at a ratio of 5:1, 83%removal was observed As the ratio decreased from 5:1 to 2:1, no increase inremoval of bromoform was observed Thus, the optimum ratio must bemaintained to achieve maximum degradation

Both H2O2 and Fe2+ are able to scavenge hydroxyl radicals generatedthrough Fenton’s reagent If any one of them is not present at the optimumdosage, either H2O2 or Fe2+ will be able to scavenge hydroxyl radicals andreduce its availability to the substrate Oxidative destruction of THMs wasfound to have slower kinetics, because THMs are saturated aliphatics andhave only one C–H bond Thus, oxidation of THMs is predominantly due

to hydrogen abstraction, which has low kinetic rates Because the order rate constant for hydrogen abstraction by hydroxyl radicals is 4 × 108

second-(M–1 s–1), a higher concentration of hydroxyl radicals is required, whichfurther implies that higher concentrations of Fe2+ and hydrogen peroxide areneeded Hydroxyl radicals are also scavenged by organic compound as well

FIGURE 6.8

4 5

6

-log[DCE Initial Concentration (M)]

as chloride and bromide ions, which can be illustrated by the followingreactions:

CHBr3 + OH•→ •CBr3 + H2O (6.92)

Fe2+ + •CBr3 + H+→ Fe3+ + CHBr3 (6.93)The above recombination reactions require the presence of either hydrogenradicals (H•) or both H+ and e– As the experiments were carried out in acidicconditions (pH = 3.5), electron transfer was possible because Fenton’s chem-istry does not generate hydrogen radicals (Huang et al., 1993)

The oxidation rates for bromoform were slower than the oxidation rates

of unsaturated chlorinated aliphatic compounds, including the TCE Becausethe hydroxylation rate constant of TCE is 109 M–1 s–1 and the hydrogenabstraction of bromoform is 1.1 × 108 M–1 s–1, aromatics and alkenes reactmore rapidly by hydroxyl addition to double bonds than does the morekinetically difficult hydrogen atom abstraction No oxidative destruction ofchloroform by Fenton’s reagent was experimentally observed; an explana-tion for this is that both H2O2 and Fe2+ have rate constants about one mag-nitude higher with respect to hydroxyl radicals than chloroform

Tang and Tassos (1997) reported that the oxidative degradation of THMsdecreases as the number of chlorine atoms present in the substrate moleculeincreases The relationship between removal rate and number of chlorineatoms was shown to be linear This phenomenon is due to the fact that thebromine substituents are better leaving groups than chlorine substituents(Solomons, 1988) Another consideration is electronegativity and bondenergy Sharp (1990) derived a relation between the bond energy betweenatoms A and B and the electronegativity as:

DAB = 0.5 (DAA + DBB) + 23 (XA – XB)2 (6.94)

where DAB, DAA, and DBB are the bond energies between A and B, A and A,and B and B, respectively Bond energy decreases as electronegativitydecreases Thus, ease of dehalogenation of an organic compound is directlyproportional to the bond energy between the carbon and halogen atoms.Because the bond energies for C–Cl and C–Br bonds are 95 kcal/mol and 67kcal/mol, respectively, brominated compounds are more easily oxidizedthan those containing proportionately more chlorine (Tang and Tassos, 1997)

Hydroxymethanesulfonic acid (HMSA) is a complex formed from dehyde and S(IV) It has been detected in atmospheric liquids (i.e., rain andsnow) The complex has high resistance to oxidation by oxygen as well asferric ions and oxygen Martin et al (1989) first studied the oxidation ofHMSA Graedel et al (1986) proposed that Fenton-type reactions are possible

formal-in atmospheric liquid water

Martin et al (1989) studied the oxidation of HMSA by Fenton’s reagentand investigated the decomposition of both hydrogen peroxide and HMSA.They determined an estimate of the absolute rate of reaction between HMSAand hydroxyl radicals The decomposition of hydrogen peroxide follows thefirst-order kinetics and can be described as follows:

–d(H2O2)/dt = k(Fe2+)(H2O2) (6.95)

where k is 0.044 M–1 s–1 at pH 2 and temperature 25°C

The actual rate of oxidation by free radicals was established by subtractingthe rate of formation of from the decomposition of HMSA Experimen-tal results showed good agreement with the first-order rate of decomposition

of HMSA Doubly ionized HMSA decomposes at a higher rate compared tosingly ionized HMSA The rate levels off until second ionization is complete,which would occur at high pH Similar experiments were performed foracetaldehyde–bisulfite complex HESA Acetaldehyde complex HESA wasnot as effective as HMSA in preventing S(IV) from being oxidized Fenton’sreagent studies were carried out at the pH levels of 1, 2, 3, and 4 Results

SO42

showed that the oxidation is a first-order process with respect to differentinitial Fe2+ concentrations and different pH at 10–1 M H2O2 and 10–2 M HMSA.

At higher concentrations of iron ion, a slight fall off was observed, whichwas attributed to less hydrogen peroxide and thus fewer hydroxyl radicals

in the system Similar experiments were carried out at various initial gen peroxide concentrations, and the oxidation was seen to be of first order;however, at larger concentrations it deviated from first order due to thesmaller amount of Fe2+ The maximum rate has been observed at pH 3.5 Anempirical rate approximating the oxidation kinetics of HMSA at pH 1 to 3

hydro-is shown below:

–d(HMSA)/(HMSA)dt = k(Fe2+)(H+)–1(H2O2)2/3 (6.96)

with k = 1.4 ± 0.2 × 10–3 (l/mol)2/3 s–1; thus, if 10–5 M H2O2 and 10–6 M Fe are

present at pH 3, the oxidation rate was 2.3 × 10–4 % h–1, which representsvery low kinetics

A study of oxidation of HMSA was done relative to pinacol to estimatethe absolute rate of oxidation of HMSA with OH radicals in solution Pinacolwas oxidized to acetone in Fenton’s oxidation Anbar and Neta (1967)reported reaction rates of OH radicals with pinacol and acetone of 3.2 × 108

M–1 s–1 and 4.3 × 107 M–1 s–1, respectively Table 6.1 presents the oxidationrates of pinacol (10–2 M) and pinacol–HMSA (10–2 M each) The concentra-

tions of the reactants were Fe2+ = 10–4 M, H2O2 = 0.1 M, and pH 2 Oxidation

rates for each molecule were different in separated and mixed reactions, asthe steady-state concentration of free radicals depends on the chemistry oforganic substrates in solution HMSA is more reactive than pinacol by afactor of 3.9 ± 0.8 If the absolute rate of reaction of pinacol with OH radicalswas calculated to be 3.2 × 108 M–1 s–1, then:

and k = 1.25 ± 0.25 × 109 M–1 s–1 This rate constant suggested that HMSAmay be consumed fairly rapidly in tropospheric clouds, as OH from othersources should be abundant to give a large reaction rate

Phenolic wastes are one of the most prevalent forms of chemical pollutants

in industry today The major sources of phenolic waste are insulation glass manufacturing, petroleum refineries, textile mills, steel making, ply-wood, hardboard production, manufacture of organic chemicals, paintstripping, and wood preservatives Eisenhauer (1964) first studied oxidation

fiber-of phenolic wastes with Fenton’s reagent It has been demonstrated that theoxidation of phenol involves the intermediate formation of catechol andhydroquinone (Merz and Waters, 1949; Stein and Weiss, 1951; Wieland and

Franke, 1928) Catechol can be oxidized in high yield to muconic acid byhydrogen peroxide and ferrous salt (Pospisil and Ettel, 1957) Thus, Eisen-hauer showed that phenol oxidation proceeds according to Equation (6.98).

(6.98)

Series of experiments for phenol concentrations of 50 ppm, at pH 3 andambient temperature of 10°C, have demonstrated that oxygen plays a majorrole in reaction by increasing the reaction rate and driving the reaction tocompletion Optimum results were obtained when the reaction was carriedout using 1 mol of ferrous salt and 3 mol of hydrogen peroxide per mole ofphenol Optimum results were obtained in a pH range of 3 to 4, as shown

in Figure 6.9

The oxidation of a number of substituted phenolic wastes by Fenton’sreagent was also studied by Eisenhauer (1964) It was observed that thegreater the degree of substitution, the slower the rate of reaction, especially

when substituents were ortho and para directing When all available positions

were blocked, as in the case of pentachlorophenol, no reaction occurred.Halophenols were rapidly oxidized and exhibited the following graduallydecreasing order: Cl > Br > I, which is attributed to the change in electrone-

gativity of the halogen groups Phenols containing meta-directing groups

such as carboxyl and nitro groups were very rapidly oxidized by Fenton’sreagent, and methyl-substituted phenols (cresols) were more resistant to

FIGURE 6.9

Effect of pH on oxidation of phenol with Fenton’s reagent (From Eisenhauer, H.R.,

J Water Pollut Contr Fed., 36, 1116-1128, 1964 With permission.)

192 Physicochemical Treatment of Hazardous Wastes

oxidation Effluents from a refinery, steel plants, and insulation plants havebeen examined A ratio of H2O2/Fe2+ of 9:1 has been able to completelyoxidize the phenol Steel plant effluent contained some cyanide, which inter-fered with the reaction; also, 16 mol of hydrogen peroxide per mole of phenoldecreased the phenol level to 35% For insulation plants, Fenton’s reagentwas also successful in reducing phenol and chemical oxygen demand (COD) Barbeni et al (1987) studied 18 different phenolic compounds The reactionwas started at pH 5 to 6, because pH decreased in the system due to gener-ation of protons by Fenton’s reagent Table 6.4 indicates that all the mono-substituted phenols were readily oxidized The effect of oxidation of di-substituted phenols was dependent on character and position of ring sub-stituents Dichlorophenols were more readily oxidized than electron-donat-ing substituents of dimethyl phenols Because Fenton’s reagent is a free-radical mechanism, any substituent that increases the electron density of thering will slow the reaction So, methyl phenol, with an electron-donatinggroup, will oxidize slowly

It was recommended that catalyst concentration levels from 10 to 20 g/giron ion should be used for phenolic compounds no greater than 2000 µg/l.Higher phenol concentrations require at least 100:1 phenol:iron ratios, andoptimum results occur when phenol solution is initially at a pH between 5and 6

Barbeni et al (1987) also compared the half-lives of different chlorophenols.They demonstrated that the appearance of chloride ion is independent ofthe disappearance of parent organic compounds during Fenton’s oxidation

TABLE 6.4

Oxidation of Phenolic Compound, 3:1 Hydrogen Peroxide:Phenol Mole Ratio

Phenolic Compound % Oxidized (1 hr, no Fe 2+ ) % Oxidized (1 hr, with Fe 2+ )

of 2-chlorophenol, and it took 30 minutes to oxidize 50 mg/L of aqueouschlorophenol (CP) Among the members of the groups of monochlorophe-

nols, meta chloro substitution induces more rapid degradation than either ortho or para substitution The formation of Cl– was slower than the disap-pearance of 3,4-DCP, which suggested that chloroaliphatic intermediate(s)may be formed after the opening of the benzene ring, and the time lagbetween appearance of Cl– ion and disappearance of phenol depends on thenumber of chloro substituents on the aromatic ring Experiments utilizingvarious ferrous concentrations have shown that increasing ferrous concen-tration in solution for a fixed hydrogen peroxide at 5 × 10–3 M increases the

rate of decomposition proportionally Similar experiments were also carriedout with ferric ions It was observed that when equal molar amount of ferrousand ferric ions were used in the same reaction mixture, the reaction half time

(t/2) was ~20 min, whereas it was ~50 min if only ferrous ions were used.

Barbeni et al (1987) indicated that the OH• formed as a result of reactionbetween Fe2+ and H2O2 are good electrophiles This electrophilic interaction

of OH• with the aromatic ring is favored by the presence of an donating group (OH)

electron-H(OH)ClC6H3H + OH• = (OH)ClC•

OH(OH)ClC6H3H(OH) = (OH)ClC6H3OH + H+ (6.100)Thus, the degradative oxidation of chlorophenols proceeds by a hydroxy-lated species (Equation 6.99 and Equation 6.100), followed by ring opening

to yield aldehydes and ultimate degradation of CO2 and Cl– It was suggestedthat the first step is the formation of radical cation by acid-catalyzed dehy-dration of radicals formed due to the interaction of OH• with chlorophenols Oxygen acts as radical scavenger and forms ; this species suc-cessively reacts with either Fe2+ or Fe3+ to regenerate H2O2 and Fe2+, as shown

in Equation (6.101) to Equation (6.103)

(6.101)

Fe2+ + → HO2• Fe3+ + HO2– → H2O2 (6.102)

Fe3+ + HO2• → Fe2+ + H+ + O2 (6.103)Under batch and semibatch conditions, Potter and Roth (1993) examinedthe oxidation kinetics of three monochlorophenol isomers and five of the six

HO•2 HO•2

dichlorophenol isomers The pH was maintained at 3.5, and the ultimateoxidation to CO2, H2O, and HCl was not reached; however, the extent ofoxidation could be estimated from the measured parameters The mechanism

of oxidation of these species is important in determining if the kinetics isfeasible The proposed reaction pathway suggested by Potter and Roth is

two possible pathways for mineralization of p-chlorophenol, as shown in

Figure 6.10

While Metelitsa (1971) proposed mineralization by substitution ofhydroxyl radicals, Pieken and Kozarich (1989) proposed mineralization byelimination (Figure 6.11) The experiments conducted by Potter and Roth(1993) showed agreement with previous work suggesting that monochlo-rophenols tend to mineralize to a greater extent than do dichlorophenols(Barbeni et al., 1987) They showed that mineralization may be favored asthe initial concentration of organic species decreases, and mineralizationoccurs primarily through other intermediates

As discussed earlier, the effects of the meta, para, and ortho positions of

chlorine on the dechlorination kinetics of monochlorophenols, nols, and trichlorophenols during Fenton oxidation were evaluated by com-paring the rate constants of the kinetic model (Tang and Huang, 1995) Thisstudy proposed a pseudo first-order steady state with respect to organicconcentration The proposed reaction pathways considered that the hydroxylradicals would attack unoccupied sites of the aromatic ring

Pentachlorophenol, a widely used wood preservative, is considered to bemoderately biorefractory with a biodegradation rate constant of 3 × 10–12 L/

cell/hr, a log K ow of 5.01, and a vapor pressure of 1.1 × 10–4 mmHg at 20°C

Watts et al (1990) carried out completely mixed batch tests by treating

penta-chlorophenol-contaminated soils with Fenton’s reagent Mineralization ofpentachlorophenol (PCP) was studied in commercially available silica sandand two natural soils by removal of parent compound and total organiccarbon with corresponding stoichiometric recovery of chloride The solubleiron concentration decreased over the first 3 hr of treatment, and the con-centration remained relatively constant thereafter A possible mechanism foriron precipitation was proposed as follows:

Fe2+ + 1/2O2 + 2OH–→ γ-FeOOH + H2O (6.104)which has a reaction half-life of 25 min at neutral pH (Sung and Morgan,1980) Watts et al (1990) proposed zero-, first-, and second-order kinetics

models but concluded that first-order was the best with r2 > 0.9 for plots

of the natural logarithm of concentration as a function of time Althoughthe reactions occurring in Fenton’s system are complex, the empirical fit

of experimental data to the first-order model has provided the most

FIGURE 6.10

Reaction pathways for oxidation of phenol (From Potter, F.J and Roth, J.A., J Hazardous Waste

Hazardous Mater., 10(2), 151, 1993 With permission.)

FIGURE 6.11

Two possible pathways for degradation of p-cholorphenol (From Potter, F.J and Roth, J.A.,

J Hazardous Waste Hazardous Mater., 10(2), 151, 1993 With permission.)