Phsicochemical Treatment of Hazardous Wastes - Chapter 6 pps

6

Fenton’s Reagent

6.1 Introduction

In 1881, Fenton published a brief description of the powerful oxidizing
properties of a mixture of hydrogen peroxide and ferrous salts. This mixture
became known as

Fenton’s reagent

, and the reaction has become known as
the

Fenton’s reaction

. Initially, Fenton applied this reaction to oxidize organic
acids such as formic, glycolic, lactic, tartronic, malic, saccharic, mucic, glyc-
eric, benzoic, picric, dihydroxytartaric, dihydroxymaleic, and acetylenedi-
carboxylic (Fenton, 1900). In the absence of ferrous salt, the degradation of
hydrogen peroxide proceeds at very slow rates, with little or no oxidation
of the organic acids (Fenton, 1899, 1900). In addition, Cross et al. (1900)
further confirmed that ferrous salts significantly enhance the kinetics of
hydrogen peroxide decomposition. Goldhammer (1927) investigated the
effect of Fenton’s reagent on phenols and found that for each equivalent of
Fe

2+

three equivalents of H

2

O

2

were decomposed. They also noted that in
concentrated hydrogen peroxide solutions each mole of Fe

2+

decomposed 24
equivalents 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 concluded
that Fenton’s reaction can be expressed as a series of chain reactions with
reaction pathways dependent on the concentration of the reactants. The
study disproved the original theory of Fenton’s reaction, which suggested
that 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 was
increased at lower pH so the propagation cycle was extended before termi-
nation. The concentration of free hydroxyl radicals was determined to be
directly proportional to the concentration of hydrogen peroxide.
Baxendale and Wilson (1957) reported that in an oxygen-free environment
Fenton’s reagent initiates very rapid polymerization of methyl acrylate,
methacrylic acid, methyl methacrylate, acrylonitrile, and styrene, and the

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166

Physicochemical Treatment of Hazardous Wastes

reaction is a function of the concentration of hydroxyl radicals. In the pres-
ence of oxygen, no polymerization occurs. Barb et al. (1951) conducted an
extensive investigation of Fenton’s reagent chemistry. When [H

2

O

2

]/[Fe

2+

]
ratios are low, the reaction rate is second order and stoichiometry is
2[Fe

2+

]

≅

[H

2

O

2

]; however, in the presence of polymerizable vinyl compound
the reaction remains second order but the stoichiometry changes to
[Fe

2+

]

≅

[H

2

O

2

]. Thus, they concluded that polymerization of vinyl compounds
occurs and results in a polymer with terminal hydroxyl groups. An inhibition

effect of hydroxyl radicals due to the higher concentration of hydrogen
peroxide was also suggested. To explain this mechanism, it was proposed
that hydroxyl radicals react with hydrogen peroxide to form hydrogen diox-
ide. This process decreases the hydroxyl radicals generated by the reaction
between ferrous iron and hydrogen peroxide. In addition, Barb et al. (1951)
suggested that hydrogen dioxide is not a strong oxidizing agent capable of
breaking the bonds of vinyl compound or oxidizing other organics.
Merz and Waters (1949) showed that oxidation of organic compounds by
Fenton’s reagent could proceed by chain as well as non-chain mechanisms,
which was later confirmed by Ingles (1972). Kremer (1962) studied the effect
of ferric ions on hydrogen peroxide decomposition for Fenton’s reagent. It
was confirmed that once ferric ions are produced the ferric–ferric system is
catalytic in nature, which accounts for relatively constant concentration of
ferrous ion in solutions.
In the late 1970s, two major theories were considered: the free radical
mechanism by Walling and Cleary (1977) and complex formation by Kremer
and Stein (1977). Walling proposed that Fenton’s oxidation predominantly
takes place by the free-radical mechanism. On the other hand, Kremer pro-
posed that complexation between the iron and the organic molecules has a
significant role and thus concluded that both mechanisms occur simulta-
neously. In the late 1980s a simultaneous effort was made to apply Fenton’s
reagent to the field of environmental science. Various contaminants were
studied in the laboratory to determine the optimum conditions. Practical
applications 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 a
reactor for treatment of organic contaminants.

6.2 Kinetic Models

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

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Fenton’s Reagent

167
proposed that iron catalyzed this reaction. For example, a small amount of
iron is sufficient to determine oxidation of an unlimited amount of tartaric
acid. In tartaric acid, two atoms of hydrogen are removed from a molecule
of acid, resulting in the production of dihydroxymaleic acid. Among com-
mon oxidants such as chlorine, potassium permanganate, atmospheric oxy-
gen, and electrolysis, the most effective oxidizing agent is hydrogen
peroxide. Fenton’s work was extended to alcohols (Fenton, 1899) and other
organic acids (1900). Attempts to identify the intermediates and products of
several organic acids and alcohols were made without success.

6.2.1 Chain Reaction Mechanism by Merz and Waters

Goldschmidt and Pauncz (1933) suggested that Fenton’s reaction is a chain
reaction involving the same reactive intermediates occurring during catalytic
decomposition of H

2

O

2

rather than via formation of peroxides of iron:
2H

2

O

2

= 2H

2

O + O

2

(6.1)
It was also shown that the ratio of oxidized alcohol to oxidized Fe

2+

could
be greater then one. Baxendale and Wilson (1957) showed that hydroxyl
radical initiating the chain polymerization of olefins by hydrogen peroxide

was the same process as the rapid oxidation of glycolic acid. Merz and Waters
(1947) confirmed that simple water-soluble alcohols are oxidized rapidly by
Fenton’s reagent. The primary alcohols are oxidized to aldehydes, which are
further oxidized at comparable rates by exactly the same mechanism. Merz
and Waters proposed a mechanism of chain oxidation of alcohols and alde-
hydes by sodium persulfate, hydrogen peroxide, and an excess of ferrous
salt as follows:
1. Chain initiation:
Fe

2+

+ H

2

O

2

= Fe

3+

+ OH

–

+ OH

•

(6.2)
2. Chain propagation:
RCH

2

OH + OH

•

= R–CHOH

•

+ H–OH (reversible) (6.3)
R–CHOH

•

+ HO–OH = R–CHO + OH

•

+ H

2

O (6.4)

3. Chain ending at low substrate concentration:
Fe

2+

+ OH

•

= Fe

3+

+ OH

–

(6.5)
4. Chain ending at high substrate (alcohol) concentration:
2R–CHOH• = R–CHO + R–CH

2

OH (disproportionation) (6.6)

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168

Physicochemical Treatment of Hazardous Wastes

In 1949, Merz and Waters determined the values for the ratio of rate
constants

k

2

/

k

3

that indicated which particular radical reduced hydrogen
peroxide. 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 is
comprised of substrates that react by non-chain processes — in this case, the
oxidation is caused by the hydroxyl radical, and considerable loss of
hydroxyl radical occurs. For the first group, the reaction rate can be expressed
by Equation (6.7):
d[H

2

O

2

]/d[RH] = 1 +

k

2

[Fe

2+

]/

k

3

[RH] (6.7)
For non-chain reactions, the kinetic rates are described by Equation (6.8):
d[H

2

O

2

]/d[RH] = 2 +

k

2

[Fe

2+

]/

k

3

[RH] (6.8)
The values for the ratio of rate constants

k

2

/

k

3

can be determined from the
intercept of their graphs. The results will suggest which particular radical
reduced hydrogen peroxide.

6.2.2 Redox Formulation by Barb et al.

Barb et al. (1951) gave a redox formulation that involves the following
reaction sequence:
(6.9)
(6.10)
(6.11)



(6.12)



(6.13)



(6.14)
Fe H O Fe HO H
3
22
2
2
•+++
+=++
k
1
Fe H O Fe OH OH
2

22
3–•++
+=++
k
2
HO H O HO HO
•
22 2 2
•
+=+
k
3
HO Fe Fe H O
2
•3 2
2
+=++
+++
k
4
HO Fe Fe OH
2
•2 3
2
–
+=+
++
k
5
OH Fe Fe OH

•2 3 –
+=+
++
k
6

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Fenton’s Reagent

169
where

k

1

and

k

2

showed inverse [H

+

] dependence.

6.2.3 Complex Mechanism by Kremer and Stein

The following scheme was presented by Kremer and Stein (1959), and further
elaborated by Kremer (1963):
(6.15)
(6.16)
FeO

3+

+ H

2

O

2

= Fe

3+

+ H

2

O

+ O

2

(6.17)
Let

C

1

= [H

+

] and

C

2

= [FeO

3+

],

k

a

and

k

d

showed inverse [H

+

] dependence
and

k

b

>>

k

a

>>

k

c

,

C

1

could be taken as a low concentration intermediate to
a good approximation
[

C

1

] =

K

[H

2

O

2

][Fe

3+

],

K

=

k

a

/

k

b

(6.18)
[Fe

3+

]

t



= [

C

2

] + [Fe

2+

] (6.19)
–d[H

2

O
2
]/dt = k
c
K[Fe
3+
]
t
[H
2
O
2
] + (k
d
– k
c

K)[C
2
][H
2
O
2
] (6.20)
d[O
2
]/dt = k
d
[C
2
][H
2
O
2
] (6.21)
d[C
2
]/dt = k
c
K[Fe
3+
]
t
[H
2
O
2

] – (k
d
+ k
c
K)[C
2
][H
2
O
2
] (6.22)
[C
2
] rises continually during the reaction, approaching a saturation value
of k
c
K[Fe
3+
]
t
/(k
c
K + k
d
), and –d[H
2
O
2
]/dt is always greater than twice d[O
2

]/
dt. At the end of the reaction, some hydrogen peroxide will be stored as
C
2
, and less than 0.5 mol of O
2
will be liberated per mole of H
2
O
2
decom-
posed.
6.2.4 Walling’s Modified Kinetic Model
Walling and Kato (1971) modified the reaction mechanism proposed by Merz
and Waters as follows:
Fe
2+
+ H
2
O
2
= Fe
3+
+ OH
–
+ OH
•
, k
1
= 76 (6.23)

Fe H O FeOOH H
3
22
2+++
++
⇔
k
k
b
a
FeOOH HO FeO
2–3++
=+
k
c
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170 Physicochemical Treatment of Hazardous Wastes
OH
•
+ Fe
2+
= Fe
3+
+ OH
–
, k
2
= 3 × 10
8

(6.24)
(6.25)
(6.26)
, k
3
= 10
7
–10
10
(6.27)
(6.28)
(6.29)
(6.30)
where k is in L/mol/s, taken from the literature. The reaction conditions
were chosen to minimize the competing processes as follows:
HO
•
+ H
2
O
2
= H
2
O + HO
2
, k = (1.2–4.5) × 10
7
(6.31)
2HO
•

= H
2
O
2
, k = 5.3 × 10
9
(6.32)
Thus, the stoichiometry is:
R = 2ar (1 – R) + b (6.33)
where R = ∆[Fe
2+
]/2∆[H
2
O
2
], a = k
2
/Σk
3
, r = [Fe
2+
]/2[RH], and b = (k
3j
+ 2k
3k
)/
2Σk
3
. This mechanism is referred to as the free-radical mechanism.
6.2.5 Ingles’ Approach

In 1972, Ingles reported his studies of Fenton’s reagent using redox titration.
He found evidence in support of Kremer’s complex mechanism theory and
concluded that, when suitable complexes are formed, substrates are not
oxidized by free radical; rather, electron transfer processes might be
OH R H H O R
•
2
•
+= +
i
i
i
k
3
OH R H H O R
•
2
•
+= +
j
j
j
k
3
OH R H H O R
•
2
•
+=+
k

k
k
3
RFeFeproduct
3
4
2
i
k
• += +
++
2R product (dimer)
•
5
j
k
=
RFe
H
Fe R H
2
6
3
kk
k
• += +
+
+
+
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Fenton’s Reagent 171
involved. Fenton’s reaction scheme was modified by Ingles for the case when
substrate is present in large amounts in the form of substrate/iron-peroxide
complexes. Ingles suggested that electron transfer occurs within this com-
plex.
(6.34)
All substrates were considered to compete as ligands in iron complexes
and 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 hydroxyl
radicals. Equation (6.37) and Equation (6.38) involve ionic mechanisms:
(6.35)
(6.36)
(6.37)
(6.38)
6.2.6 Transition State Approach by Tang and Huang
6.2.6.1 Competitive Method
When modeling oxidation kinetics of chlorophenols by Fenton’s reagent,
elementary rate constants are critical to obtain quantitative stoichiometric
data in terms of optimal dosages for H
2
O
2
and Fe
2+
to achieve a given removal
efficiency. If an elementary reaction rate constant for a given compound is
not available in the literature, another method can be used to determine it

experimentally. For example, the rate constants of 2,4-dichlorophenol (DCP)
and 2,4,6-trichlorophenol (TCP) were determined by an alternative method
by Tang and Huang (1996a). The equation used to calculate the rate constants
is as follows:
R
I
Fe
II
OOH
R
I
Fe
III
O
+
OH
-
+
.
R
2
R
2
R
I
Fe
II
OOH
.
R

I
Fe
III
O
+
+
OH
-
R
I
Fe
II
OOH
+
R
I
Fe
II
O
+
OH
-
+
R
I
Fe
II
OOH
R
I

Fe
III
O
+
OH
-
+
+
R
2
R
2
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172 Physicochemical Treatment of Hazardous Wastes

(6.39)
where:
= rate constant between any organic compound and hydroxyl
radical.
= rate constant between reference compound and hydroxyl rad-
ical.
[S] = concentration of the substrate at any time.
[S
0
] = initial concentration of the substrate.
[R] = concentration of the reference compound at any time.
[R
0
] = initial concentration of the reference compound (2-chlorophenol).

In their work, the reference compound is 2-chlorophenol, with a rate constant
of 8.2 × 10
9
M
–1
s
–1
. Either 2,4-DCP or 2,4,6-TCP was mixed with 2-chlorophe-
nol in a reactor, separately. Then, H
2
O
2
was mixed with the organic com-
pound and the pH was adjusted to 3.5. The organic concentrations were
measured by gas chromatography (GC) before and after Fe
2+
was added.
The results are shown in Figure 6.1. According to the slopes of the straight
line, the rate constants between hydroxyl radicals and 2,4-DCP and 2,4,6-
TCP can be determined as 7.22 × 10
9
M
–1
s
–1
and 6.27 × 10
9
M
–1
s

–1
, respectively.
The hydroxylation rate constants for 2,4-DCP and 2,4,6-TCP are clearly
smaller than that for 2-chlorophenol; therefore, increasing chlorine content
on the aromatic ring decreases the reactivity of the chlorinated phenols
toward hydroxyl radical attack.
6.2.6.2 Dechlorination Kinetic Model
6.2.6.2.1 Pseudo First-Order Kinetic Model
When an excess of H
2
O
2
and Fe
2+
is added at constant concentrations to the
system, a steady-state concentration of hydroxyl radical can be assumed.
The concentration of both H
2
O
2
and Fe
2+
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: k
1

is the pseudo first-order rate constant of oxidation.
k
2

Intermediates +
•
OH = chloride ion + CO
2
+ other products (6.41)
kk
HO ,S
0
0
HO ,R
S]/[S
R]/[R
••
=
ln([ ])
ln([ ])
k
HO ,S
•
k
HO ,R
•
k
1
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Fenton’s Reagent 173
where k
2
is the pseudo first-order rate constant of dechlorination. The deg-
radation kinetics can be modeled as the following:
d(CP)/dt = –k
1
(CP) (6.42)
d(I)/dt = k
1
(CP) – k
2
(I) (6.43)
d(Cl
–
)/dt = k
2
(I) (6.44)
where CP is the concentration of chlorinated phenols at any time t; I is the
concentration of intermediates formed at any time t; and Cl
–
is the concen-
tration of chloride ion. The integrated form of the above equation is:
(CP)/(CP)
0
= exp(–k
1
t) (6.45)
(Cl)/(CP)
0

= 1 + [k
1
exp(–k
2
t) – k
2
exp(–k
1
t)]/(k
2
– k
1
) (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 for
the oxidation and dechlorination of 2,4,6-TCP.
FIGURE 6.1
The competitive oxidation kinetics of different chlorinated phenols at optimal ratio of H
2
O
2
to
Fe
2+
of 25 and optimal pH of 3.5. Experimental conditions: (2-CP) = (2,4-DCP) = (2,4,6-TCP) =
5 × 10
–4
M; H

2
O
2
= 5 × 10
–3
M; Fe(ClO4)
2
= 2 × 10
–4
M; pH = 3; ionic strength = 0.05 M as Na
2
SO
4
.
In[(Reference)/(Reference)
0
]
-In[(Substrate)/Substrate)
0
]
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174 Physicochemical Treatment of Hazardous Wastes
It is important to note that both H
2
O
2
and Fe
2+
have to be overdosed to

maintain a steady-state concentration of hydroxyl radical and to obtain a
satisfactory approximation of the mathematical model with the experimen-
tal data. When H
2
O
2
and Fe
2+
concentrations are 5 × 10
–3
M and 2 × 10
–4
M,
respectively, the relative rate constants of 2-chlorophenol (2-CP) and 2,4,6-
TCP with respect to 2,4-DCP can be calculated. The oxidation and dechlo-
rination constants of 2,4-DCP were found to be 0.995 1/min (k
1
) and 0.092
1/min (k
2
), as reported in a previous study (Tang and Huang, 1996). For
comparison, 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 math-
ematical model, k
1
= 0.150 1/min and k
2
= 0.032 1/min). Experimental conditions: 2,4,6-TCP =

5 × 10
–4
M; H
2
O
2
= 5 × 10
–3
M, Fe(ClO4)
2
= 2 × 10
–4
M; pH = 3.5; ionic strength = 0.05 M as N
2
SO
4
.
TABLE 6.1
Kinetic Rate Constants of Chlorinated Phenols by Fenton’s Reagent
2-CP 2,4-DCP 2,4,6-TCP
Elementary rate constants (M
–1
s
–1
) 8.2 × 10
9
7.2 × 10
9
6.3 × 10
9

Measured oxidation constants
(1/min)
1.666 0.995 0.15
Measured dechlorination constants
(1/min)
1.386 0.092 0.032
0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
0.9
1
051015 20 25 30
Time (minutes)
Normalized 2,4,6-TCP or Cl- concentration
Cl- concentration detected, conc. calculated by model
2.4.6-TCP measured, conc.calculated according to model
Cl-
2,4,6-TCP
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Fenton’s Reagent 175
To evaluate the effect of the number of chlorines on the degradation rate
constants of different chlorophenols, Table 6.2 shows the rate constants of
elementary, oxidation, and dechlorination for the ratios of k

2-CP
/k
2,4-DCP
and
k
2,4,6-TCP
/k
2,4-DCP
. The relative rate constants are plotted against the number of
sites unoccupied by chlorine atoms on the chlorinated phenols in Figure
6.3.A linear correlation between the rate constants and the number of sites
available is found with a standard deviation of 0.132. Clearly, the more
chlorine atoms the aromatic rings contain, the fewer sites are available for
hydroxyl 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/k
2,4-DCP
)
2-CP 2,4-DCP 2,4,6-TCP
Elementary rate constants 1.14 1 0.88
Observed oxidation constants 1.67 1 0.15
Observed dechlorination constants 15.07 1 0.35
FIGURE 6.3
The correlation between oxidation constants and sites available on aromatic ring for hydroxyl
radical attack. Experimental conditions: 2-CP = 2,4-DCP = 2,4,6-TCP = 5 × 10
–4
M; H
2
O

2
= 5 ×
10
–3
M; Fe(ClO4)
2
= 2 × 10
–4
M; pH = 3.5; ionic strength = 0.05 M as N
2
SO
4
.
y=0.755X - 1.328
0
0.5
1
1.5
2
234
Number of Sites Available
Relative Rate Constants to K
2,4-DCP
relative rate constants measured for different chlorophenols
correlation of the rate constants of chlorophenol
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176 Physicochemical Treatment of Hazardous Wastes
predicting the rate constants of tetra-chlorophenol and penta-chlorophenol
due to steric hindrance. Figure 6.3 indicates that the oxidation rate decreases

with the increasing degree of chlorine content in the following order (in
terms of both elementary rate constants and the observed pseudo first-order
rate constants): 2-CP > 2,4-DCP > 2,4,6-TCP.
6.2.6.2.2 Dechlorination Kinetic Model Using Transition State Theory
When both H
2
O
2
and Fe
2+
are not overdosed, the concentrations of both
H
2
O
2
and Fe
2+
will change. The pseudo first-order kinetic model developed
above does not apply. In order to quantitatively model the effect of H
2
O
2
and Fe
2+
on the dechlorination kinetics by Fenton’s reagent, the dechlorina-
tion kinetic model is developed as follows. First, hydroxyl radicals are gen-
erated by H
2
O
2

decomposition by Fe
2+
:
Fe
2+
+ H
2
O
2
→ Fe
3+
+
•
OH + OH
–
k
i
= 51 M
–1
s
–1
(6.47)
where k
i
is the initiation rate constant of hydroxyl radical generation. Then,
the following termination reactions occur simultaneously in the reactor:
2
•
OH → H
2

O
2
k
t1
= 5.3 × 10
9
M
–1
s
–1
(6.48)
Fe
2+
+
•
OH → Fe
3+
+ OH
–
k
t2
= 3 × 10
8
M
–1
s
–1
(6.49)
•
HO + H

2
O
2
→ H
2
O + HO
2
k
t3
= 2.7 × 10
7
M
–1
s
–1
(6.50)
Tang and Huang (1996a) reached two conclusions. First, hydroxyl radicals
will attack unoccupied sites of the aromatic ring; second, chlorine atoms will
be released from the chlorinated aliphatic intermediates instead of the aro-
matic ring. When chlorinated phenols are present in the system, the follow-
ing reaction mechanisms can be assumed:
(6.51)
(6.52)
(6.53)
where k
f
is the rate constant for formation of activated complex; k
b
is the rate
constant for decomposition of activated complex after hydroxyl radical

CP OH (CP OH)
•*
+
→
←
−
k
k
b
f
(CP OH) chlorinated aliphatic intermediates
*
−→
k
m
Chlorinated aliphatic compounds OH CO Cl H O
•
2
–
2
+→ ++
k
p
TX69272_C06.fm Page 176 Friday, November 14, 2003 2:09 PM
© 2004 by CRC Press LLC
Fenton’s Reagent 177
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, and
chlorinated intermediates, their concentrations are extremely low at the
steady state; therefore, a pseudo first-order steady state can be assumed for
the kinetic modeling. As a result, the steady-state concentration of the acti-
vated complex can be obtained by setting the change of its concentration to
zero:
(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 equal
zero:

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

(6.59)
d(C)
d
( OH) *(CP) (C) (C) 0
*

**
t
=
k

k

k
=
fbm
⋅ ––
(C)
( OH) (CP)
*
=
k
*
k
+
k
f
bm
⋅
d(I)
d
(C (I)( OH) 0
*
t
=
k

k
=
mp
)–
•
(I)
(
C
)
( OH)
*
=
k
k
m
p
⋅
d( OH)
d
(
H
O
)(
Fe
)(OH)(CP) (C )
2
2
2+
*

⋅
⋅
t
=
kk
+(
kk
)
if bm
––

kk k
=
tt t
–– –
1
2
2
2+
3
2
2
( OH) ( OH)(
Fe
)(OH)(
H
O
)0⋅⋅ ⋅
iff
bm

bm
kk
+
k
(
kk
)
(
k
+
k
)
(
H
O
)(
Fe
)(OH)(CP) ( OH)(CP)
2
2
2+
–
–
⋅⋅

kk k
=
tt t
–– –
1

2
2
2+
3
2
2
(OH
)
( OH)(
Fe
)(OH)(
H
O
)0⋅⋅ ⋅
TX69272_C06.fm Page 177 Friday, November 14, 2003 2:09 PM
© 2004 by CRC Press LLC
178 Physicochemical Treatment of Hazardous Wastes
Simplifying this equation, we get:

(6.60)
For simplicity, we assume that the following expression containing rate
constants is also a constant k:
(6.61)
Then, the final form of hydroxyl radical concentration can be expressed as
follows:
(6.62)
The rate equation for dechlorination of chlorinated phenols by Fenton’s
reagent can be expressed as follows:
(6.63)
so the final expression for dechlorination is:

(6.64)
Because Fenton oxidation of chlorophenols should follow the same mech-
anism, the activated complex in the transition state should have a similar
structure. Therefore, Equation (6.64) can be applied to mono-, di-, and trichlo-
rophenols. It is not certain, however, that it can be applied to tetra- and penta-
chlorophenols due to steric hindrance; therefore, when the above general
equation is applied to chlorophenols, the equation becomes:
(6.65)
if
m
bm
kk
(
k
)
(
k
+
k
)
(
H
O
)(
Fe
)2 (OH)(CP)
2
2
2+
– ⋅

tt
kk
=
2
2+
3
2
2
( OH)(
Fe
)(OH)(
H
O
)
−
−
⋅⋅0
k =
kk
k
+
k
fm
bm
2
()
( OH)
(
H
O

)(
Fe
)
(CP) (
Fe
)(
H
O
)
2
2
2+
2
2+
3
2
2
⋅ =
k
k +
k
+
k
i
tt
r=
t
=
k
=

k
k
k
+
k
=
k
Cl
pm
f
bm
–
d(
Cl
)
d
(I)( OH)
(CP)( OH)
2
(CP)( OH)
–
⋅
⋅
⋅
d(
Cl
)
d
2(CP) (
H

O
)(
Fe
)
CP (
Fe
)(
H
O
)
–
2
2
2+
2
2+
3
2
2
t
=
k/
k
k( ) +
k
+
k
i
tt
1

2 (CP) (
H
O
)
(
H
O
)
(CP)
1
Fe
2
(
H
O
)
(CP)
3
2
2
0
2
2
0
0
2+
2
2
2
0

0
r
=
k +
k
k
k
+
k
k
k
0
t
i
t
i
0
[]
TX69272_C06.fm Page 178 Friday, November 14, 2003 2:09 PM
© 2004 by CRC Press LLC
Fenton’s Reagent 179
The k can be obtained using the initial rate method by overdosing H
2
O
2
. For
example, at the overdosed H
2
O
2

concentration of 5 × 10
–3
M and constant
pH of 3.5, the k can be obtained by plotting the 1/r
0
vs. 1/Fe
2+
during the
dechlorination 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
be calculated as 7 × 10
5
(1/s).
When Fe
2+
is overdosed, the k value can be obtained through the following
equation:
(6.66)
At the overdosed Fe
2+
concentration of 5 × 10
–3
M and constant pH of 3.5,
the 1/r
0
vs. 1/H
2
O
2
during the dechlorination of 2,3,4-TCP is shown in Figure

6.5. From the slope of 54.1 (s), the k can be calculated as 7.3 × 10
5
(1/s), which
is about the same as when H
2
O
2
was overdosed.
Using the developed model, the k values for 2-CP, 3-CP, and 4-CP are 1.12
× 10
7
, 1.004 × 10
9
, and 1.005 × 10
8
(1/s), respectively; therefore, the dechlori-
nation constants for monochlorophenols follow a decreasing order: 3-CP >
4-CP > 2-CP. Because chloride ion can be released only after the rupture of
the aromatic ring, the faster the hydroxylation of the parent compounds, the
faster the dechlorination process should be. Therefore, the above order can
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 of
hydroxyl radical attack is expected:
FIGURE 6.4
The plot of reciprocal initial dechlorination rate vs. reciprocal concentration of Fe
2+
during 2,3,4-
TCP oxidation. Experimental conditions: H
2

O
2
= 5 × 10
–3
M; 2,3,4-TCP = 5 × 10
–4
M; pH = 3.5;
ionic strength = 0.05 M as N
2
SO
4
.
y=7.21x + 1.94 * 10
4
0
20000
40000
60000
80000
100000
120000
140000
160000
180000
200000
0 5000 10000 15000 20000
1/(Fe
2+
) (1/M)
1/r

0
(s/M)
Initial rate of dechlorination measured for 2,3,4-TCP
Correlation model
1
2 (TCP) (
Fe
)
(
Fe
)
(TCP)
1
H
O
2
(
Fe
)
(TCP)
0
0
0
0
0
0
0
2
2+
2+

2
2
3
2+
r
=
k +
k
k
k
+
k
k
k
t
i
t
i
[]
TX69272_C06.fm Page 179 Friday, November 14, 2003 2:09 PM
© 2004 by CRC Press LLC
180 Physicochemical Treatment of Hazardous Wastes
(6.67)
In Relation (6.67), the solid arrow presents a stronger directory effect by
the 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, inter-
mediates with higher degrees of oxidation are expected to be produced in
oxidation 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
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 is
the 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 expe-
rience more steric strain than meta or para monochlorophenols.
Figure 6.6 presents the plots of dechlorination kinetic rate constants vs.
H
2
O
2
/(CP) ratio during the oxidation of 2-MCP, 2,4-DCP, and 2,4,6-TCP at
a constant H
2
O
2
/Fe
2+
ratio of 2.5 and optimal pH of 3.5. When the ratio of
H
2
O
2
/DCP increases, the dechlorination rate constants increase. Further-
more, the difference between dechlorination rate constants becomes more
FIGURE 6.5
The plot of reciprocal initial dechlorination rate vs. reciprocal concentration of H

2
O
2
during
2,3,4-TCP oxidation. Experimental conditions: Fe
2
= 5 × 10
–3
M; 2,3,4-TCP = 5 × 10
–4
M; pH =
3.5; ionic strength = 0.05 M as N
2
SO
4
.
y=54.1x + 1380
0
50000
100000
150000
200000
0 500 1000 1500 2000
1/(H
2
O
2
) (1/M)
1/r
0

(s/M)
Initial rate of dechlorination measured for 2,3,4,-TCP
Correlation model
OH
OH
OH
Cl
Cl
Cl
>
>
TX69272_C06.fm Page 180 Friday, November 14, 2003 2:09 PM
© 2004 by CRC Press LLC
Fenton’s Reagent 181
and more apparent when the H
2
O
2
/DCP ratio increases from 1 to 20. For a
constant H
2
O
2
/CP, the H
2
O
2
/Fe
2+
ratio is constant. Therefore, the steady-

state concentration of hydroxyl radicals is constant. However, as the ratio of
H
2
O
2
/DCP increases, more hydroxyl radicals are available to react with CP.
Because chloride ion is released through numerous hydroxylation and
hydrogen abstraction steps, more hydroxyl radicals are available for a DCP
molecule, and the faster the dechlorination kinetics will be. The effect of the
H
2
O
2
/CP ratio on the dechlorination constants can be clearly seen in Figure
6.6. The dechlorination rate follows this decreasing order: k
TCP
> k
DCP
> k
MCP
.
At the optimal H
2
O
2
/Fe
2+
ratio, the amount of H
2
O

2
available for 1 mol of
chlorinated phenols affects dechlorination rate constants significantly, as
shown in Figure 6.9. When the ratio of H
2
O
2
/CP is 20, the dechlorination
constants 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 are
1:3:6. At a constant H
2
O
2
/CP ratio, the maximal capacity of Fenton’s reagent
is constant, because the number of available sites not occupied by chlorine
atoms is the major factor responsible for dechlorination. If we assume that
each 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 gives
1:2.6:6 as a relative dechlorination rate constant. The theoretical prediction
agrees fairly well with the experimental result of 1:3:6. This is the reason
why no linear correlation could be found between dechlorination constants
and the nonchlorinated sites available. All the sites unoccupied by chlorine
FIGURE 6.6
Dechlorination rate constants for 2-CP, 2,4-DCP, and 2,4,6-TCP vs. H
2
O
2
/organic concentrations

for 2-CP, 2,4-DCP, and 2,4,6-TCP. Experimental conditions: H
2
O
2
= 5 × 10
–3
M; Fe
2+
= 2 × 10
–3
M;
pH = 3.5; ionic strength = 0.05 M as N
2
SO
4
.
0
0.05
0.1
0.15
0.2
0.25
0.3
0.35
02468101214161820
(H
2
O
2
)/(CP

)
Rate Constants of Dechlorination K
Cl-
(1/s)
2,46-TCP 2,4-DCP 2-MCP
TX69272_C06.fm Page 181 Friday, November 14, 2003 2:09 PM
© 2004 by CRC Press LLC
182 Physicochemical Treatment of Hazardous Wastes
atoms on the aromatic ring of the chlorinated phenols have the same reac-
tivity toward hydroxyl radicals when chlorine atoms occupy the 2, 4, and 6
positions. 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
> 3,5-DCP > 2,3-DCP > 2,6-DCP > 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; dechlorina-
tion has the following order in terms of decreasing rate constants: 2,4,6-TCP
> 2,4,5-TCP > 2,3,4-TCP. This order suggests that the steric hindrance
becomes a determining factor in preventing hydroxyl radical attack on the
unoccupied sites of aromatic rings. For example, 2,4,6-TCP has much larger
space than 2,4,5-TCP due to separation of the chlorine atoms on the aromatic
ring. 2,4,5-TCP, in turn, has a larger space than does 2,3,4-TCP in which
chlorine atoms can locate closely to one another. Therefore, 2,4,6-TCP is
affected the least by steric hindrance; 2,4,5-TCP is subjected to an average
steric 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, if
the 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 the

initial dechlorination rate is independent of the organic concentration:
k(CP) >> [k
t2
(Fe
2+
) + k
t3
(H
2
O
2
)] (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 that
the limiting step in the oxidation of chlorinated phenols is the generation of
hydroxyl radicals through Fenton’s reagent.
(6.69)
TABLE 6.3

Dechlorination Rate Constants during
Oxidation of Dichlorophenols
Dichlorophenol
Rate Constants
(×10
7
1/s)
2,4-Dichlorophenol 2.28
2,3-Dichlorophenol 8.61
2,5-Dichlorophenol 11.90
2,6-Dichlorophenol 13.46
3,5-Dichlorophenol 13.82
3,4-Dichlorophenol 14.12
r=
k
i
0
(
H
O
)
(
Fe
2
2
0
2+
0
)
TX69272_C06.fm Page 182 Friday, November 14, 2003 2:09 PM

© 2004 by CRC Press LLC
Fenton’s Reagent 183
6.2.6.2.3 Oxidation Model of Unsaturated Aliphatic Compounds
The same transition complex approach and steady-state assumptions were
used to develop the kinetic model of unsaturated chlorinated aliphatic com-
pounds such as trichloroethylene (TCE). The model reflects the effects of
H
2
O
2
, Fe
2+
, and organic compounds on the oxidation kinetics as follows:
(6.70)
where:
(6.71)

The above equation agrees with the experimental observation that oxidation
kinetics is the pseudo first-order reaction with respect to the concentration
of substrates. The rate constant of k
observed
can be experimentally determined
FIGURE 6.7
Dechlorination initial rate vs. different organic concentrations. Experimental conditions: H
2
O
2
= 5 × 10
–3
M; Fe

2+
= 2 × 10
–3
M; pH = 3.5; ionic strength = 0.05 M as N
2
SO
4
.
4
4.1
4.2
4.3
4.4
4.5
4.6
4.7
4.8
4.9
5
2 2.5 3 3.5 4
-log (CP)
-log(r
0
(M/S))
2-MCP 2,4-DCP 2,4,6-TCP average r
o
r=
k
*
k

k
+
k
(S)
s
observed
n
i
tt
(
H
O
)(
Fe
)
(
Fe
)(
H
O
)
2
2
2+
2
2+
3
2
2







observed
pf
bp
k
=
kk
k
+
k

TX69272_C06.fm Page 183 Friday, November 14, 2003 2:09 PM
© 2004 by CRC Press LLC
184 Physicochemical Treatment of Hazardous Wastes
by keeping two variables constant and varying one variable among the
concentrations of H
2
O
2
, Fe
2+
, and substrate. When the logarithm of both sides
of Equation (6.70) is taken, the following linear relationship can be obtained:
(6.72)
Let X be equal to the following:
(6.73)

The equation can be written as:
(6.74)
Using the experimentally determined constants such as k
i
, k
+2
, and k
+3
, the
kinetic models for dichloroethylene (DCE), TCE, and tetra-CE can be
expressed as follows:

(6.75)

(6.76)
(6.77)
It is important to point out that the above treatment is not applicable to
different oxidation mechanisms. For example, the kinetic data obtained from
dichloroethane oxidation does not give a straight line between log(r
0
) and
log(X), because the oxidation pathway of dichloroethane is different from
that in oxidizing chlorinated ethylenes. For example, dichloroethane is one
log log log log(r ) = (
k
) +n*
k
k
+
k

+ (S )
observed
i
tt
0 0
(
H
O
)(
Fe
)
(
Fe
)(
H
O
)
2
2
0
2+
0
2
2+
0
3
2
2
0







X =
k
k
+
k

i
tt
(
H
O
)(
Fe
)
(
Fe
)(
H
O
)
2
2
0
2+
0

2
2+
0
3
2
2
0
log( ) log log logr = n* X + (
k
) + (S )
observed
00
r=
t
= *
k
k
+
k
*
s
i
tt
–
d(DCE)
d
0.21
(
H
O

)(
Fe
)
(
Fe
)(
H
O
)
(DCE)
0.439
2
2
2+
2
2+
3
2
2






r=
t
= *
k
k

+
k
*
s
i
tt
–
d(TCE)
d
19.95
(
H
O
)(
Fe
)
(
Fe
)(
H
O
)
(TCE)
0.261
2
2
2+
2
2+
3

2
2






r=
(Tetra - )
t
= *
k
k
+
k
* (Tetra - )
s
i
tt
–
dCE
d
3.71
(
H
O
)(
Fe
)

(
Fe
)(
H
O
)
CE
0.333
2
2
2+
2
2+
3
2
2






TX69272_C06.fm Page 184 Friday, November 14, 2003 2:09 PM
© 2004 by CRC Press LLC
Fenton’s Reagent 185
of the saturated chlorinated aliphatic compounds, and the first oxidative
step is hydrogen abstraction instead of hydroxylation (Walling, 1975). The
unsaturated bond formed after hydrogen abstraction is then attacked by the
addition of hydroxyl radicals. As a result, hydrogen abstraction must precede
hydroxylation in the oxidation of dichloroethane.

In order to obtain the optimal ratio of H
2
O
2
to Fe
2+
, Equation 6.76 can be
differentiated with respect to H
2
O
2
, assuming that Fe
2+
is the optimal con-
centration of (Fe
2+
)
opt
. We set:
(6.78)
The optimal concentration of H
2
O
2
can be expressed as follows by solving
the above equation.
(6.79)
The optimal concentration of Fe
2+
can be derived from the same mathemat-

ical approach. We set the derivative of r
s
with respect to Fe
2+
concentration
equal to zero:
(6.80)
Then, the optimal concentration of Fe
2+
is:
(6.81)
When Equation (6.79) is divided by Equation (6.81), the optimal ratio of
H
2
O
2
/Fe
2+
can be obtained:
(6.82)
Therefore, the final form of the optimal ratio between H
2
O
2
and Fe
2+
can be
expressed as:
(6.83)
d

d(
H
O
)
0
2
2
s
r
=
(
H
O
)
(
Fe
)
2
2
2
2+
3
opt
t2
opt
t
=
k
(S)
k

–
d
d(
Fe
)
0
2+
s
r
=
(
Fe
)
(
H
O
)
2+
3
2
2
2
opt
t
opt
t
=
k
(S)
2

k
–
(
H
O
)
(
Fe
)
(
Fe
)
(
H
O
)
2
2
2+
2
3
2+
2
2
=
(
k
)
(
k

)
*
t
2
t
2
opt






HO
Fe
22
2
2
3
()
()











=
+
opt
t
t
k
k
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© 2004 by CRC Press LLC
186 Physicochemical Treatment of Hazardous Wastes
Substituting the numerical values into the above equation, we obtain the
optimal ratio of H
2
O
2
to Fe
2+
as the following:
(6.84)
Experimental data indicate that the optimal ratio of H
2
O
2
to Fe
2+
is about
5 to 11. To investigate the effect of a low H
2
O

2
/Fe
2+
ratio, the Fe
2+
concen-
tration was overdosed at a constant value of 10
–3
M. At a high H
2
O
2
/Fe
2+
ratio, where H
2
O
2
was overdosed at 10
–2
M, the extrapolated maximum initial
rate of 1.8 mM/min also occurs at an H
2
O
2
/Fe
2+
ratio of 11. This value
reasonably agrees with the theoretical value of 11 as the optimal ratio of
H

2
O
2
/Fe
2+
. At a constant H
2
O
2
/Fe
2+
, the H
2
O
2
concentration required for iso-
percentage release of chloride ion in dichloroethylene is shown in Figure
6.8.Tang and Huang (1997) concluded that the amount of H
2
O
2
required for
a specific percentage removal of the organic compounds depends upon the
initial organic concentration to be oxidized. This is also true for the total
percentage of chloride ion released at different initial organic concentrations.
The typical percentage removal of organic compounds and percentage
release of chloride ion have been studied at 100, 70, 50, 40, 30, 20, 10, and
1%. The amount of H
2
O

2
required to achieve a certain percentage removal
follows the order of TCE < tetra-CE < DCE << DCEA (dichloroethane) at a
Fe
2+
concentration of 10
–3
M. However, the amount of chloride ion detected
at an H
2
O
2
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 ali-
phatic compounds.
6.3 Oxidation of Organic Compounds
6.3.1 Trihalomethanes
Trihalomethanes (THMs) are priority pollutants listed by the U.S. Environ-
mental Protection Agency (EPA). They are recalcitrant in nature, thus their
destruction is difficult. The most commonly encountered THMs in drinking
water threatening human health are chloroform, bromodichloromethane,
dibromochloromethane, and bromoform. Tang and Tassos (1997) studied the
oxidation kinetics and mechanisms of these four THMs.
The effect of the ratio of H
2
O
2

to Fe
2+
on oxidation kinetics, the oxidation
kinetics 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 to
oxidize of the four THMs. Bromoform concentrations used in the study of
Fenton’s reagent ratio and oxidation kinetics were 49.2, 98.3, and 295 µg/L.
As the ratio of H
2
O
2
to Fe
2+
increases, the removal efficiency increases with
(
H
O
)
(
Fe
)
3*
10
2.7*
10
11
2
2
2+
2

3
8
7






opt
t
t
=
k
k
= =
TX69272_C06.fm Page 186 Friday, November 14, 2003 2:09 PM
© 2004 by CRC Press LLC
Fenton’s Reagent 187
the initial concentration of bromoform. This may indicate that the hydroxyl
radical has a preference toward organic compounds, resulting in proportion-
ately less scavenging effect by H
2
O
2
:Fe
2+
and MeOH. For a higher H
2
O

2
:Fe
2+
ratio of 10:1, the amount of bromoform removal appeared to show depen-
dence on initial bromoform concentration. As the initial organic substrate
concentration increases, less scavenging of OH
•
occurs. At an H
2
O
2
:Fe
2+
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 in
removal of bromoform was observed. Thus, the optimum ratio must be
maintained to achieve maximum degradation.
Both H
2
O
2
and Fe
2+
are able to scavenge hydroxyl radicals generated
through Fenton’s reagent. If any one of them is not present at the optimum
dosage, either H
2
O
2

or Fe
2+
will be able to scavenge hydroxyl radicals and
reduce its availability to the substrate. Oxidative destruction of THMs was
found to have slower kinetics, because THMs are saturated aliphatics and
have only one C–H bond. Thus, oxidation of THMs is predominantly due
to hydrogen abstraction, which has low kinetic rates. Because the second-
order rate constant for hydrogen abstraction by hydroxyl radicals is 4 × 10
8
(M
–1
s
–1
), a higher concentration of hydroxyl radicals is required, which
further implies that higher concentrations of Fe
2+
and hydrogen peroxide are
needed. Hydroxyl radicals are also scavenged by organic compound as well
FIGURE 6.8
H
2
O
2
concentration required for iso-percentage release of chloride ion in dichloroethylene (DCE)
oxidation at optimal conditions. Experimental conditions: H
2
O
2
/Fe
2+

= 5; pH = 3.5; ionic strength
= 0.05 M as Na
2
SO
4
.
2
3
4
5
23456
-log[DCE Initial Concentration (M)]
-log[H2O2 Required (M)]
1% 10% 50% 100%
TX69272_C06.fm Page 187 Friday, November 14, 2003 2:09 PM
© 2004 by CRC Press LLC
188 Physicochemical Treatment of Hazardous Wastes
as chloride and bromide ions, which can be illustrated by the following
reactions:
OH
•
+ RH → H
2
O + R
•
(6.85)
OH + R → ROH
•
(6.86)
CHBr

3
+ OH
•
→
•
CBr
3
+ H
2
O (6.87)
•
CBr
3
+ OH
•
→ HO-CBr
3
(6.88)
HO = CRr
3
+ OH
•
→ O = CBr
2
+ H
2
O + Br
–
(6.89)
Br

–
+ OH
•
→ BrOH
–
(6.90)
The inhibitory effect of chloride and bromide ions is not significant unless
free ion concentration exceeds a range from 1 × 10
–2
to 5 × 10
–2
M. Dehalo-
genation is believed to be a slower process than oxidation of the parent
compound, and the proposed general reaction is as follows:
THM + OH
•
→ halogenated intermediates → X
–
+ CO
2
+ H
2
O + end products
(6.91)
Fluctuations in bromoform concentration were observed even after a tripli-
cate run. This was explained by recombination of the bromoform radical as
follows:
CHBr
3
+ OH

•
→
•
CBr
3
+ H
2
O (6.92)
Fe
2+
+
•
CBr
3
+ H
+
→ Fe
3+
+ CHBr
3
(6.93)
The above recombination reactions require the presence of either hydrogen
radicals (H
•
) or both H
+
and e
–
. As the experiments were carried out in acidic
conditions (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. Because
the hydroxylation rate constant of TCE is 10
9
M
–1
s
–1
and the hydrogen
abstraction of bromoform is 1.1 × 10
8
M
–1
s
–1
, aromatics and alkenes react
more rapidly by hydroxyl addition to double bonds than does the more
kinetically difficult hydrogen atom abstraction. No oxidative destruction of
chloroform by Fenton’s reagent was experimentally observed; an explana-
tion for this is that both H
2
O
2
and Fe
2+
have rate constants about one mag-
nitude higher with respect to hydroxyl radicals than chloroform.
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Fenton’s Reagent 189
Tang and Tassos (1997) reported that the oxidative degradation of THMs
decreases as the number of chlorine atoms present in the substrate molecule
increases. The relationship between removal rate and number of chlorine
atoms was shown to be linear. This phenomenon is due to the fact that the
bromine substituents are better leaving groups than chlorine substituents
(Solomons, 1988). Another consideration is electronegativity and bond
energy. Sharp (1990) derived a relation between the bond energy between
atoms A and B and the electronegativity as:
D
AB
= 0.5 (D
AA
+ D
BB
) + 23 (X
A
– X
B
)
2
(6.94)
where D
AB
, D
AA
, and D
BB
are the bond energies between A and B, A and A,
and B and B, respectively. Bond energy decreases as electronegativity

decreases. Thus, ease of dehalogenation of an organic compound is directly
proportional 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 67
kcal/mol, respectively, brominated compounds are more easily oxidized
than those containing proportionately more chlorine (Tang and Tassos, 1997).
6.3.2 Hydroxymethanesulfonic Acid
Hydroxymethanesulfonic acid (HMSA) is a complex formed from formal-
dehyde and S(IV). It has been detected in atmospheric liquids (i.e., rain and
snow). The complex has high resistance to oxidation by oxygen as well as
ferric ions and oxygen. Martin et al. (1989) first studied the oxidation of
HMSA. Graedel et al. (1986) proposed that Fenton-type reactions are possible
in atmospheric liquid water.
Martin et al. (1989) studied the oxidation of HMSA by Fenton’s reagent
and investigated the decomposition of both hydrogen peroxide and HMSA.
They determined an estimate of the absolute rate of reaction between HMSA
and hydroxyl radicals. The decomposition of hydrogen peroxide follows the
first-order kinetics and can be described as follows:
–d(H
2
O
2
)/dt = k(Fe
2+
)(H
2
O
2
) (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 subtracting
the 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 to
singly ionized HMSA. The rate levels off until second ionization is complete,
which would occur at high pH. Similar experiments were performed for
acetaldehyde–bisulfite complex HESA. Acetaldehyde complex HESA was
not as effective as HMSA in preventing S(IV) from being oxidized. Fenton’s
reagent studies were carried out at the pH levels of 1, 2, 3, and 4. Results
SO
4
2–
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© 2004 by CRC Press LLC