Ozone Reaction Kinetics for Water and Wastewater Systems - Chapter 2 potx

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2

Reactions of Ozone
in Water

Due to its electronic configuration, ozone has different reactions in water. These fall
into three categories:
•Oxidation–reduction reactions
•Dipolar cycloaddition reactions
• Electrophilic substitution reactions
A possible fourth way of reaction could be some sort of nucleophilic addition,
although it has only been checked in nonaqueous systems.

1

In some cases, free radicals are formed from these reactions. These free radicals
propagate themselves through mechanisms of elementary steps to yield hydroxyl
radicals. These hydroxyl radicals are extremely reactive with any organic (and some
inorganic) matter present in water.

2

For this reason, ozone reactions in water can be
classified as direct and indirect reactions. The direct reactions are the true ozone
reactions, that is, the reactions the molecule of ozone undergoes with any other type
of chemical species (molecular products, free radicals, etc.). The indirect reactions
are those between the hydroxyl radical, formed from the decomposition of ozone,
or from other direct reactions of ozone, with compounds present in water. It can be

said that direct ozone reactions are the initiation step leading to indirect reactions.

2.1 OXIDATION–REDUCTION REACTIONS

Redox reactions are characterized through the transfer of electrons from one species
(reductor) to another one (oxidant).

3

The oxidizing or reducing character of any
chemical species is given by the standard redox potential. Ozone presents one of
the highest standards of redox potentials,

4

only lower to those of the fluorine atom,
oxygen atom, and hydroxyl radical (see Table 2.1). Because of its high standard
redox potential, the ozone molecule presents a high capacity to react with numerous
compounds by means of this reaction type. This reactivity is particularly important
in the case of some inorganic species such as Fe

2+

or

Γ

. However, in most of these
reactions there is no explicit transfer of electron, but rather an oxygen transfer from
the ozone molecule to the other compound. Examples of explicit electron transfer

reactions are scarce but the reactions between ozone and the hydroperoxide ion and
the superoxide ion radical could be catalogued in this group

6

:
(2.1)OHO O HO
32 3 2
+→•+•
−−

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(2.2)
In most of the cases, however, one oxygen atom is transferred as, for example, in
the reaction with Fe

2+

:
(2.3)
Nonetheless, in all these reactions, some atom of the inorganic species goes to
a higher valence state, that is, it looses electrons, so that these reactions could
theoretically be catalogued as oxidation–reduction reactions since, in an implicit
way, there is an electron transfer. The reaction of ozone with nitrite is an example
of this. The two half reactions are
(2.4)
(2.5)
The standard redox potential allows checking the possibility that ozone reacts
through redox reactions with a given compound. Main electron ion half reactions

of ozone in water are reactions (2.5) and (2.6):
(2.6)

TABLE 2.1
Standard Redox Potential of Some Oxidant Species

5

Oxidant Species E

o

, Volts Relative Potential of Ozone

Fluorine 3.06 1.48
Hydroxyl radical 2.80 1.35
Atomic oxygen 2.42 1.17
Ozone 2.07 1.00
Hydrogen peroxide 1.77 0.85
Hydroperoxide radical 1.70 0.82
Permanganate 1.67 0.81
Chlorine dioxide 1.50 0.72
Hypochlorous acid 1.49 0.72
Chlorine 1.36 0.66
Bromine 1.09 0.53
Hydrogen peroxide 0.87 0.42
Iodine 0.54 0.26
Oxygen 0.40 0.19
OO O O
32 3 2

+•→•+
−−
OFe FeO O
3
2
22
+→+
++
NO H O e NO H
22 3
22
−−−+
+−→+
OH e OHO
322
22++→+
+−
OHOe O OH E v
o
32 2
2124++→+ =
−−
.

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From these data the importance of pH on ozone

redox

reactions can be deduced.
Detailed information of standard redox potential of different substances can be
obtained elsewhere.

3,4

2.2 CYCLOADDITION REACTIONS

Addition reactions are those resulting from the combination of two molecules to
yield another one.

7

One of the molecules usually presents atoms-sharing more than
two electrons (i.e., unsaturated compounds such as olefinic compounds with a carbon
double bond) and the other one presents an electrophilic character. These unsaturated
compounds present

π

electrons that in a lesser extent keep bonded the carbon atoms
of the double bond. These

π

electrons are quite available to electrophilic compounds.
It can also be said that an addition reaction develops between one base compound
(a compound with

π

electrons) and an acid compound (an electrophilic compound).
As a general rule, the following scheme would correspond to an addition reaction:
(2.7)
In practice, there could be different types of addition reactions such as those
between ozone and any olefinic compound. In this case, the reaction follows the
mechanism of Criegge

8

that constitutes an example of cycloaddition reaction. The
mechanism of Criegge develops through three steps as shown in Figure 2.1. In the
first step, a very unstable five-member ring or primary ozonide is formed.

9

This
breaks up, in a second step, to give a zwitterion. In the third step, this zwitterion
reacts in a different way, depending on the solvent where the reaction develops, on
experimental conditions, and on the nature of the olefinic compound. Thus, in a
neutral solvent, it decomposes to yield another ozonide, some peroxide or ketone,
and polymer substances as shown in Figure 2.2. When the reaction is in a partici-
pating solvent (i.e., a protonic or nuclephilic solvent) some oxy-hydroperoxide
species is generated (Figure 2.3). Finally, a third possibility is the so-called abnormal
ozonolysis that could develop both in participating and nonparticipating solvents.
In this way, some ketone, aldehyde, or carboxylic acids can be formed (Figure 2.4).
The cycloaddition reaction, then, leads to the break up of both

σ

and

π

bonds of the
olefinic compound while the basic addition reaction (2.7) leads only to the break up
of the

π

bond. Compounds with different double bonds (C=N or C=O) do not react
with ozone through this type of reaction.

10,11

This is not the case of aromatic
compounds that could also reacts with ozone through 1,3-cycloaddition reactions
leading to the break up of the aromatic ring. However, in these cases, the cycload-
dition reaction also is less probable than the electrophilic attack of one terminal
oxygen of the ozone molecule on any nucleophilic center of the aromatic compound.
The reason of this is due to the stability of the aromatic ring because of the resonance.
Notice that the cycloaddition reaction leads to the break up of the aromatic ring,
then to the loss of aromaticity, while the electrophilic reaction (see later) retains the
aromatic ring.
−=−+ →− − −CC XY XC CY

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FIGURE 2.1

Criegee mechanism.

FIGURE 2.2

Decomposition ways of primary ozonide in an inert solvent.

FIGURE 2.3

Decomposition ways of primary ozonide in a participating solvent.
C C C C
C
C
Different ways of reaction
(see Figures 2.2 to 2.4)
III
I
II
C
C
C
C
2
C
+
2
CC
x
( )
CC
CC

C
-PH
C
OOH
P
-P = OH
C
OOH
OH
-P
=
-N
H
C
OOH
N
H
-P
=
-C
O
O
C
OOH
O-C
-P = -O
OOH
O
O
O

C
O

©2004 CRC Press LLC

2.3 ELECTROPHILIC SUBSTITUTION REACTIONS

In these reactions, one electrophilic agent (such as ozone) attacks one nucleophilic
position of the organic molecule (i.e., an aromatic compound), giving rise to the
substitution of one part (i.e., atom, functional group, etc.) of the molecule.

7

This
type of reaction is the base of the ozonation of aromatic compounds such as phenols
as shown later. Aromatic compounds are prone to undergo electrophilic substitution
reactions rather than cycloaddition reactions because of the stability of the aromatic
ring. For example, the benzene molecule is strongly stabilized by the resonance
phenomena. The benzene molecule can be represented by different electronic struc-
tures that constitute the benzene hybrid. The difference of stability between individ-
ual structures and the hybrid is the energy of resonance. In the case of benzene, the
individual structure is the cyclohexatriene, and the resonance energy is 36 kcal, that
is, the energy difference between those of the cyclohexatriene and the benzene
hybrid. This is the reason of the aromatic properties: the higher the resonance energy,
the stronger the aromatic properties. The reactions of aromatic compounds depend
on these aromatic properties. Thus, after the electrophilic substitution, the aromatic
properties are still valid, and the resulting molecules present the aromatic stability.
This situation is lost when cycloaddition takes place.
In a general way, an aromatic substitution reaction develops in two steps as
shown in Figure 2.5 for the case of benzene and one electrophilic agent YZ. In the

first step, a carbocation (C

6

H

5
+

HY) is formed and, in the second, a proton is taken
due to the action of a base compound.

FIGURE 2.4

Examples of abnormal ozonolysis.

FIGURE 2.5

Basic steps of the aromatic electrophilic substitution reaction.
C
C
C
C +C C
C
R
R
C
Ketone
H
R

C
Aldehyde
HO
R
C
Acid
+
H
+ E
Slow
H
E
H
E
+ :N
Fast
E
+ H:N

©2004 CRC Press LLC

Another important fact to consider is the presence of substituting groups in the
aromatic molecule (i.e., phenols, cresols, aromatic amines, etc.). These groups
strongly affect the reactivity of the aromatic ring with electrophilic agents. Thus,
groups such as HO–, NO

2
–

, Cl

–

, etc., activate or deactivate the aromatic ring for the
electrophilic substitution reaction. Also, depending on the nature of the substituting
group, the substitution reaction can take place in different nucleophilic points of the
aromatic ring. Thus, activating groups promote the substitution of hydrogen atoms
from their ortho and para positions with respect to these groups, while the deacti-
vating groups facilitate the substitution in the meta position. Table 2.2 shows the
effect of different substituting groups on the electrophilic reaction of the benzene
molecule. In fact, both the resulting products of the electrophilic substitution reaction
and the relative importance of the reaction rate can be predicted after considering
the nature of substituting groups. Differences in the rate of substitution reaction
should theoretically be due to differences in the slow step of the process, that is, the
formation of the carbocation: the higher the stability of the carbocation, the faster
the electrophilic substitution reaction rate. The carbocation is a hybrid of different
possible structures where the positive charge is distributed throughout the aromatic
ring, although positions ortho and para, regarding the substituting group position,
present the higher nucleophilic character. As a consequence, these positions have
the highest probability to undergo the electrophilic substitution reaction (see Figure 2.6).
Factors that affect the spreading of the positive charge are those that stabilize the
carbocation or intermediate state.
Also, the substituting group can increase or decrease the carbocation stability,
depending on the capacity to release or take electrons. From Figure 2.6, it is evident that
the stabilizing or destabilizing effect is especially important when the substituting group
is bonded to the ortho or para carbon atom with respect to the attacked nucleophilic

TABLE 2.2
Activating and Deactivating Groups of the Aromatic
Electrophilic Substitution Reaction

7

Groups Action on Reaction Importance

–OH

–

, –O

–

, –NH

2

, –NHR, –NR

2

Activation Strong
–OR, –NHCOR Activation Intermediate
–C

6

H

5

, –Alkyl Activation Weak
–NO

2

, –NR

3
+

Deactivation Strong
–C

≡

N, –CHO, –COOH Deactivation Intermediate
–F, –Cl, –Br, –I Deactivation Weak

FIGURE 2.6

Resonance forms of the hybride carbocation.
H
E
H
E
H
E
H
E

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position. Groups such as alkyl radicals or –OH activate the aromatic ring because they
tend to release electrons while groups such as –NO

2

deactivate the aromatic ring since
they attract electrons. In the first case, the carbocation is stabilized, while in the second
case it is not. For example, in the case of the ozonation of phenols, this property is
particularly important due to the strong electron donor character of the hydroxyl group.
In addition, the carbocation formed in the case of phenol is a hybrid constituted not only
by the contribution of structures I to III (see Figure 2.6) but also by a fourth structure
(see Figure 2.7) where the positive charge is on the oxygen atom. Structure IV is
especially stable since each atom (except the hydrogen atom) has completed the orbitals
(eight electrons). This carbocation is more stable than those from the electrophilic sub-
stitution in the benzene molecule (where there is no substituting group) or in the meta
position with respect to the –OH group in the molecule of phenol (Figure 2.8). In these
two cases, structure IV is not possible, then the ozonation of phenol is faster than that
of benzene and goes mainly at ortho and para positions with respect to the –OH group.
In fact, literature reports kinetic studies (see Chapters 3 and 5) of the ozonation of
aromatic compounds where the rate constant of the direct reactions between ozone and
phenol, and ozone and benzene have been found to be 2

×

10

6

and 3

M

–1

sec

–1

, respec-
tively.

12–14

It should be noticed, however, that these values correspond to pH 7 and 20˚C.
As shown later, rates of phenol ozonation are largely influenced by the pH of water
because of the dissociating character of phenols. More information on the stability of
carbocations in electrophilic substitution reactions in different aromatic structures can
be obtained from organic chemistry books.

7

In the case of the ozonation of phenol, the mechanism goes through different
electrophilic substitution and cycloaddition reactions as shown in Figure 2.9.

15–17

2.4 NUCLEOPHILIC REACTIONS

According to the resonance structures of the ozone molecule (see Figure 1.2), there
exists a negative charge on one of the terminal oxygen atoms. This fact confers, at

FIGURE 2.7

Resonance forms of the carbocation formed during the ozonation of phenol
(attack to ortho position).

FIGURE 2.8

Resonance forms of the carbocation formed during the ozonation of phenol
(attack to meta position).
H
E
H
E
OH OH
H
E
OH
H
E
OH
H
E
OH
H
E
OH

H
E
OH

©2004 CRC Press LLC

least theoretically, a nucleophilic character to the ozone molecule. Thus, ozone could
react with molecules containing electrophilic positions. These reactions belong to
the nucleophilic addition type, and molecules with double (and triple) bonds between
atoms of different electronegativity could theoretically be involved. In the case of
ozonation, the nucleophilic activity can be shown in the presence of carbonyl or
double and triple carbon nitrogen bonds.

1

Thus, the following example shows two
possible ways (nucleophilic and electrophilic) of an ozone attack on a ketone. For
example, the nucleophilic reaction of ozone on Schiff bases with carbon-nitrogen
double bonds has been reported. Figure 2.10 shows this example. It should be noted,
however, that most of the information related to the mechanism of the ozonation of
organic compounds has been obtained in an organic medium, and that there is scarce
information on this matter when water is the solvent.

2.5 INDIRECT REACTIONS OF OZONE

These reactions are due to the action of free radical species coming from the
decomposition of ozone in water. The free radical species are formed in the initiation
or propagation reactions of the mechanisms of advanced oxidation processes involv-
ing ozone and other agents, such as hydrogen peroxide or UV radiation, among
others.

18

An advanced oxidation process (AOP) is defined as that producing hydroxyl
radicals which are strong oxidant species.

2

In the ozone decomposition mechanisms,

FIGURE 2.9

General mechanism of the ozonation of phenol (AO = Abnormal ozonolysis).
HO
OH
O
O
O O
OO
HO
OH
OH
OH
HO
OH
OH
O
O
CO
2

+ H
2
O
C
C
O
HO
O
HO
H
2
O
2
O
HO
C
C
H
OH
C
O
OH
C
O
OH
C
O
OH
C
O

H
H
2
O
2
H
2
O
2
HO
C
O
OH
O
H
C
OH
O
C
C
O
C
H
H
O
C
C
O
C
OH

O
C
C
O
OH
HO
O
C
C
O
H
HO
O
C
C
O
H
H
AO
AO
AO

©2004 CRC Press LLC

the hydroxyl radical is the main responsible species of the indirect reactions. Then,
the reaction between the hydroxyl radical and compounds (that could be called
pollutants) present in water constitute the indirect reactions of ozone.
Numerous studies have been developed to clarify the mechanism of decompo-
sition of ozone in water since Weiss in 1934

19

proposed the first model. Today, the
mechanism of Staehelin, Hoigné, and Buhler (SHB)

20–23

is generally accepted that
ozone follows its decomposition in water, although when pH is high, another mech-
anism by Tomiyasu, Fukutomi, and Gordon (TFG) is also considered as the most
representative.

24

In Tables 2.3 and 2.4 both mechanisms together to values of the
rate constants of their reactions are shown.
The reactions of ozone with the hydroxyl and hydroperoxide ions can be con-
sidered as the main initiation reactions of the ozone decomposition mechanism in
water. However, other initiation reactions develop when other agents, such as UV
radiation or solid catalysts are also present. Thus, the direct photolysis of ozone that
yields hydrogen peroxide and then free radicals,

26

or the ozone adsorption and
decomposition on a catalyst surface to yield active species (in some cases hydroxyl
radicals,

27

as will be discussed in other chapters) are also examples of initiation
reactions. The reaction of ozone and the superoxide ion radical [reaction (2.2)] is
one of the main propagating reactions of the ozone decomposition mechanism.
There are also other reactions that lead to the decomposition or stabilization of
ozone in water. Thus, substances of different nature can also contribute to the appear-
ance or inhibition of free radicals. These substances are called initiators, inhibitors,
and promoters of the decomposition of ozone.

21

The initiators are those substances,
such as the hydroperoxide ion (the ionic form of hydrogen peroxide) mentioned
above, that directly react with ozone to yield the superoxide ion radical [reaction
(2.1)]. These reactions are initiation reactions. The superoxide ion radical is the key
to propagating free radical species because it rapidly reacts with ozone to yield free

FIGURE 2.10

An ozone nucleophilic substitution reaction. (From Riebel, A.H. et al., Ozo-
nation of carbon-nitrogen bonds. I. Nucleophilic attack of ozone,

J. Am. Chem. Soc

., 82,
1801–1807, 1960. With permission.)
+
OOO
C
6
H

5
CH N R
C
6
H
5
CH N R C
6
H
5
CH N R
OO
O
OO
O
OO
O
C
6
H
5
CH N R
HHH
C
6
H
5
CH N R
O
C

6
H
5
CHO N RC
6
H
5
C NHR
O
+

©2004 CRC Press LLC

radicals, such as the ozonide ion radical [reaction (2.2)] that eventually leads to the
hydroxyl radical (see Table 2.3 or 2.4). Promoters are those species that, through their
reaction with the hydroxyl radical, propagate the radical chain to yield the key free
radical: the superoxide ion radical. Examples of these substances are methanol, formic
acid, or some humic substances.

21

Of particular interest is the role of hydrogen
peroxide in the mechanism of ozone decomposition. In fact, hydrogen peroxide is
the initiating agent of ozone decomposition as proposed by Tomiyasu et al.

24

but it
also acts as promoter of ozone decomposition according to the following reactions

28

:

TABLE 2.3
Ozone Decomposition Mechanism in Pure Water According
to Staehelin, Hoigné, and Bühler

22,23



Reaction Rate constant Reaction #
Initiation Reaction

a


70

M

–1

sec

–1

(2.8)

Propagation Reactions


7.9

×

10

5

sec

–1



25

(2.9)

5

×

10

10

M

–1

sec

–1



25

(2.10)

1.6

×

10

9



M

–1

sec

–1

(2.2)

5.2

×

10

10



M

–1

sec

–1

(2.11)

3.3

×

10

2

sec

–1

(2.12)

1.1

×

10

5

sec

–1

(2.13)

2

×

10

9

M

–1

sec

–1

(2.14)

2.8

×

10

4

sec

–1

(2.15)

Termination Reactions



b

5

× 10
9
M
–1
sec
–1

25
(2.16)
b
5 × 10
9
M
–1
sec
–1

25
(2.17)
a
Later, Hoigné
6
considered reaction (2.8) should be reaction (2.18)] of Tomiyashu
et al. mechanism
24
(see Table 2.4) although the rate constant value kept the same
(70 M–1sec–1). This reaction change implies that reaction (2.1), hydrogen peroxide

equilibrium reactions (2.22) and (2.23) (see Table 2.4) and reactions between
hydrogen peroxide and the hydroxyl radical (2.27) and (2.28) also take part of the
mechanism.
b
Reaction products, H
2
O
2
and O
3
were tentatively proposed.
OOH HO O
k
i
322
1
+→•+•
−−
HO O H
k
22
1
• →•+
−+
OH HO
k
22
1
−+
′

•+  →•
OO O O
k
32 3 2
2
+•→•+
−−
OH HO
k
33
3
−+
•+  →•
HO O H
k
33
4
• →•+
−+
HO HO O
k
32
5
• →•+
OHO HO
k
34
6
+•→•
HO HO O

k
422
7
• →•+
HO HO H O O
k
T
44 223
1
2•+ • →•+
HO HO H O O O
k
T
43 2223
2
•+ • →•++
©2004 CRC Press LLC
(2.27)
(2.28)
However, as shown in Chapter 8, hydrogen peroxide can also act as indirect inhibitor
of the ozone decomposition, when its concentration is so high the ozone/hydrogen
peroxide reaction becomes mass transfer-controlled.
TABLE 2.4
Ozone Decomposition Mechanism in Pure Water at Alkaline
Conditions According to Tomiyasu, Fukutomi, and Gordon
24
Reaction Rate Constant Reaction #
Initiation Reaction
*
40 M

–1
sec
–1
(2.18)

2.2 × 10
6
M
–1
sec
–1
(2.1)
Propagation Reactions

7.9 × 10
5
sec
–1

25
(2.9)

5 × 10
10
M
–1
sec
–1

25

(2.10)

1.6 × 10
9
M
–1
sec
–1
(2.2)

20–30 M
–1
sec
–1
(2.19)

6 × 10
9
M
–1
sec
–1
(2.20)

3 × 10
9
M
–1
sec
–1

(2.21)

5 × 10
10
M
–1
sec
–1

25
(2.22)

0.25 sec
–1

25
(2.23)
Termination Reactions

2.5 × 10
9
M
–1
sec
–1
(2.24)
*
4.2 × 10
8
M

–1
sec
–1
(2.25)
*
No data was given (2.26)
*
Carbonates were assumed to be present because of alkaline conditions. In fact,
reactions (2.25) and (2.26) are not true termination reactions since the superoxide
ion radical, O
2
–
•, would propagate the radical chain. Reaction products, O
2
, CO
2
and O
2
–
• were tentatively proposed but not confirmed. Reactions (2.27) and (2.28)
(see text) have to be added to this mechanism.
OOH HO O
k
322
8
+→+
−−
OHO HO O
k
i

32 23
2
+→•+•
−−
HO O H
k
22
9
• →•+
−+
OH HO
k
22
9
−+
′
•+  →•
OO O O
k
32 3 2
2
+•→•+
−−
OHO HOOOH
k
32 2
10
−−
•+  →•++
OHO HOO

k
322
11
−−
•+ • →•+•
OHO HO O
k
32
6
+•→+•
HO H H O
k
222
12
−+
+→
HO HO H
k
22 2
12
′
−+
→+
OHO OOH
k
T
33
3
+•→+
−

HO CO OH CO
k
C
•+  →+•
=−−
33
2
CO O O CO O
k
T
33 2 22
4
−−
•+  →++•()
HO H O HO H O
kMs
H
•+  →•+
=×
−−
22
27 10
22
1
711
.
HO HO HO OH
kMs
H
•+  →•+

−
=×
−
−−
2
75 10
2
2
911
.
©2004 CRC Press LLC
Finally, inhibitors of the ozone decomposition are those species that while
reacting with the hydroxyl radical terminate the radical chain. In this group, one can
cite tert-butanol, p-chlorobenzoate ion, carbonate, and bicarbonate ions or, also,
some other humic substances.
21,27
The inhibitors are also called hydroxyl-free radical
scavengers because their presence limit or avoid the action of these radicals on the
target contaminants. For example, the presence of carbonates in natural water reduces
the efficiency of ozonation to oxidize refractory contaminants also present in the water.
Because of the importance of these three types of substances, different alternative
mechanisms to those of SHB or TFG have been proposed. Thus, it is particularly
significant that the mechanisms of ozone decomposition in the presence of carbonate
species (carbonate and bicarbonate ions), due to their usual presence in water, are
called natural inhibitors. However, carbonate ion species can not be considered as
pure inhibitors of ozone decomposition. In this case, some other reactions must be
added to the mechanisms shown in Tables 2.3 and 2.4, especially if hydrogen
peroxide is present in significant concentration. These reactions are shown in Table
2.5. As can be seen from Table 2.5, the reactions of hydroxyl radicals with carbonate
species yield the carbonate ion radical. This free radical is not inactive in many

cases. Instead, the carbonate ion radical is able to regenerate the carbonate ions by
reacting with hydrogen peroxide (see reactions in Table 2.5). Also, the carbonate
ion radical can react with some substances (i.e., phenol) and constitute another way
of oxidation.
31–33
More information on the rate constant of these reactions can be
obtained from other works.
2,31
Another case, extensively studied in the literature because of its health impact,
is the presence of bromide ion in ozonated water.
34,35
Ozone easily oxidizes bromide
ion to yield a toxic pollutant, bromate ion. As indicated in Chapter 1, environmental
agencies have established a low MCL for bromate ion in water. The reactions of
bromide–ozone processes are shown in Table 2.6. It can be seen that there are
different reactions between the species that appear in this mechanism. Formation of
bromate is highly dependent on the presence of other different substances that
consume ozone such as hydrogen peroxide or ammonia that react with hypobromous
acid to yield bromamines.
42,43
Another important aspect often considered in the ozone decomposition mecha-
nism in water is the presence of natural organic matter (NOM). Depending on the
nature of NOM, these substances can act as promoters or inhibitors of the decom-
position of ozone. For this reason, the following reaction is usually included in the
mechanism of ozone decomposition when NOM is present
44,45
:
(2.68)
or
(2.69)

Consideration of these reactions basically means that a fraction of hydroxyl radical,
while reacting with NOM, yields the superoxide ion radical and hence a fraction of NOM
is a promoter of the ozone decomposition reaction. With this reaction, it is accepted that
NOM HO O P+•→•+
−
α
2
NOM HO O P+•→•+
−
α
2
©2004 CRC Press LLC
part of the matter present in water acts as a promoter and another part acts as inhibitor.
It should be noted that the kinetic behavior of ozone in natural water is similar than in
wastewater but with the difference that direct ozone reactions and action of promoters,
initiators, and inhibitors can be multiple due to the complexity of the organic and
inorganic matrix of the wastewater. In Chapter 6 aspects related to the treatment of ozone
kinetics of wastewater are presented.
2.5.1 THE OZONE DECOMPOSITION REACTION
The fact that ozone once dissolved in water is unstable and decomposes
constitutes one of its advantages but also one drawback. Thus, on the
one hand, when ozone decomposes, free radicals, particularly the hydroxyl radical
is generated and oxidation of compounds develops in a process called
advanced oxidation (indirect reaction). On the other hand, because
TABLE 2.5
Main Reactions Involving Carbonate Species in Water
during Ozonation Processes
2,25,29–31
Reaction Rate constant Reaction #

8.5 × 10
6
M
–1
sec
–1
(2.29)

4.2 × 10
8
M
–1
sec
–1
(2.30)

2.2 sec
–1
*
(2.31)

5 × 10
10
*
(2.32)

2.25 × 10
4
*
(2.33)

5 – 10
10
*
(2.34)

500 sec
–1
*
(2.35)

5 × 10
10
M
–1
sec
–1
*
(2.36)

4.3 × 10
5
M
–1
sec
–1
(2.37)

5.6 × 10
7

M
–1
sec
–1
(2.38)

7.5 × 10
8
M
–1
sec
–1
(2.39)

6 × 10
7
M
–1
sec
–1
(2.40)

See Reference 31 (2.41)
*
For the rate constant of protonation reactions a value of 5 – 10
10

25
has been
assumed considering that these reactions are diffusion controlled. The rate

constant of the inverse reaction has been calculated from the corresponding
pK and the indicated value for the protonation reaction.
HCO HO HCO OH
k
C
33
1
−−
+•→•+
CO HO CO OH
k
C
33
2
=−−
+•→•+
HCO CO H
k
C
33
3
−=+
→+
HCO H CO
k
C
33
3
−+
′

=
+→
HCO HCO H
k
C
23 3
4
→+
−+
HCO H H CO
k
C
323
4
−+
′
+→
HCO CO H
k
C
33
5
• →•+
−+
CO H HCO
k
C
33
5
−+

′
•+  →•
CO H O HCO HO
k
C
322 3 2
6
−−
•+  →+•
CO HO CO HO
k
C
32 32
7
−− =
•+  →+•
CO O CO O
k
C
32 32
8
−− =
•+ • →+
CO O CO O
k
C
33 33
9
−− =
•+ • →+

CO B CO
k
C
33
10
−=
•+  →+Products
©2004 CRC Press LLC
TABLE 2.6
Main Reactions Involving Bromine Species during
Ozonation Processes
34,36–42
Reaction Rate constant Reaction #

50 M
–1
sec
–1
(2.42)

300 M
–1
sec
–1
(2.43)

100 M
–1
sec
–1

(2.44)

5 × 10
10
M
–1
sec
–1
(2.45)

50 sec
–1
(2.46)

10
5
M
–1
sec
–1
(2.47)

1.1 × 10
10
M
–1
sec
–1
(2.48)

3.3 × 10
7
M
–1
sec
–1
(2.49)

4.2 × 10
6
M
–1
sec
–1
(2.50)

2 × 10
9
M
–1
sec
–1
(2.51)

10
10
M
–1
sec
–1

(2.52)

2 × 10
9
M
–1
sec
–1
(2.53)

8.24 M
–1
sec
–1
(2.54)

8 × 10
7
M
–1
sec
–1
(2.55)

4.5 × 10
9
M
–1
sec
–1

(2.56)

2 × 10
9
M
–1
sec
–1
(2.57)

9.5 × 10
8
M
–1
sec
–1
(2.58)

7 × 10
4
M
–1
sec
–1
(2.59)

2 × 10
5
M
–1

sec
–1
(2.60)

4.1 × 10
9
M
–1
sec
–1
(2.61)

4.9 × 10
9
M
–1
sec
–1
(2.62)

2 × 10
9
M
–1
sec
–1
(2.63)

1.4 × 10
9

M
–1
sec
–1
(2.64)

7 × 10
9
sec
–1
(2.65)

7 × 10
8
M
–1
sec
–1
(2.66)

4.3 × 10
7
M
–1
sec
–1
(2.67)
Br O BrO O
−−
+→+

32
BrO O Br O
−−
+→+
32
2
BrO O BrO O
−−
+→+
322
BrO H HBrO
−+
+→
HBrO BrO H→+
−+
BrO O BrO O
23 32
−−
+→+
Br HO BrOH
−−
+•→•
BrOH Br HO
−−
•→+•
BrOH Br OH
−−
•→•+
Br Br Br•+ →•
−−

2
Br • +O BrO •+O
32
→
2
23
Br Br Br
−−
•→•+•
Br H O HBrO Br H
22
+→++
−+
BrO Br BrO Br
−− −
+•→•+
2
2
BrO HO BrO OH
−−
+•→•+
HBrO HO BrO H O+•→•+
2
HBrO O Br OH O+•→•++
−
22
_
HBrO H O Br H O H O+→+++
−+
22 2 2

BrO H O Br H O
−−
+→+
22 2
BrO Br Br BrO
−−
+•→+•
22
22
BrO H O BrO BrO H•+ →++
−−+
BrO HO BrO OH
22
−−
+•→•+
2
224
BrO Br O•→
Br O BrO
24 2
2→•
Br O OH BrO BrO H
24 3 2
+→++
−−−+
CO BrO CO BrO
33
−− =
•+ →+•
©2004 CRC Press LLC

of the instability, ozone can not be used in practice as final disinfectant
of water. The kinetic study of the decomposition of ozone in water is
one of the necessary steps to ascertain whether or not ozone is able
to remove some given compounds from water through direct or indirect
reactions.
The ozone decomposition rate, as can be deduced from the previous section, is
highly dependent on the nature of substances present in water. For example, Figure 2.11
presents data of decomposition of ozone in buffered distilled water at three pH
values. As deduced from Figure 2.11 (and also from the mechanisms of Tables 2.3
and 2.4), pH is also one of the main factors that influence the decomposition of
ozone in water. As a general rule, for pH < 7 this variable has a slight effect on the
ozone decomposition,
46
but at higher pH, the rate increases significantly. As example,
Figure 2.12 shows the variation of the apparent rate constant of the decomposition
of ozone with pH in a study where first-order kinetics was considered.
46
Because of the importance of hydroxyl radical oxidation, the decomposition of
ozone in water has been the subject of numerous works. Since the first work on this
matter due to Weiss,
19
numerous researchers have studied the ozone decomposition
kinetics from reaction mechanisms, based on experimental facts. The rate equations
are finally fitted to the experimental data of ozone concentration time obtained in
homogeneous ozone decomposition reactions in water. Thus, rate equations of dif-
ferent complexity with one or two terms, first, second, or three halves order with
respect to ozone, with the rate constant expressed as a function of temperature and
pH, have been deduced and checked in literature (see Table 2.7). As is observed
from Table 2.7, the reaction order changes in many cases and also the influence of
pH. From the SHB mechanism initiators, promoters, and inhibitors of the decompo-

sition of ozone are considered as the responsible substances of the kinetic differences
between rate equations reported in Table 2.7. For a natural water, application of the
SHB mechanism leads to the following ozone decomposition rate equation:
FIGURE 2.11 Decomposition of ozone in buffered distilled water at different pH. Conditions:
17˚C, buffered water with phophates, ⅙ pH 2, C
O30
.= 8.1 mgl
–1
, ● pH 7, C
O30
.= 4.2 mgl
–1
, ▫
pH 8.5, C
O30
.= 3.7 mgL
–1
.
Ozone conversion
Time, min
0
0.2
0.4
0.6
0.8
1
0 10 20 30 40 50 60
©2004 CRC Press LLC
(2.70)
where C

Mi
, C
Ii
, C
Pi
, and C
Si
are the concentrations of compounds that react directly
with ozone, initiators, promoters, and inhibitors compounds, respectively, and C
O3
the concentration of dissolved ozone.
In spite of the works of Hoigné et al.,
20–23
studies on this matter still continue
as shown in Table 2.7. The determination of the rate constant and reaction order
was usually carried out by following the decrease of the ozone concentration in
water with time in an agitated tank through homogeneous experiments (see Chapter
3). Due to the effects of the natural organic matter in water, when projecting any
study of ozonation where the decomposition of ozone is thought to be important, it
is recommended to determine the specific rate law of ozone decomposition if the
water matrix is different from that used in the studies reported previously. In fact,
there are two clear ozone decomposition periods when treating natural waters.
74
The
first one, called fast ozone demand, varies just from a few seconds to as much as
approximately 1 or 2 min. During this period an instantaneous or very fast consump-
tion of ozone takes place, and during the second or long ozone decomposition period,
ozone slowly decomposes. It is this latter period that the subject of kinetics studies:
ozone decomposition in natural waters. The initial ozone fast-demand period is due
to the presence of substances that readily react with ozone through direct reactions.

Once these substances have disappeared or their concentration become lower the
longer ozone decomposition period starts. More information on the ozone decom-
position kinetics is given in Chapter 7.
FIGURE 2.12 Variation of the apparent pseudo first-order rate constant of ozone decompo-
sition in buffered distilled water with pH. T, ºC: ⅙ 10, ∆ 20, ▫ 30, ∇ 40. (From Sotelo, J.L.
et al., Ozone decomposition in water: Kinetic Study, Ind. Eng. Chem. Res., 26,39–43, 1987.
With permission.)
pH
,
min
–1
k
d
10
–2
10
–1
1
2 4 6 8 10
−= = + +
{}
+



















∑∑
∑
∑
−
rkC kC kC kC
kC
kC
C
O
d
ODiMii
OH
Ii Ii
Pi Pi
Si Si
O33 3
31
©2004 CRC Press LLC
TABLE 2.7
Works on Aqueous Ozone Decomposition Kinetics

Reacting System and Operating Conditions Main Observations Reference #
Batch reactor, iodometric analysis
0ºC, pH 2–4, H
2
SO
4
or HNO
3
Rate equation: –r
O3
= A(a–C
O3
)
2
+ B(a–C
O3
).
Data on k determined. No influence of pH
47
(1913)
Batch reactor, iodometric analysis
0ºC, pH 5-8, phosphate buffers
–r
O3
= k10
0.36(pH 14)
C
O3

Data on k determined

48
(1933)
Data from other works at 0ºC and acid pH Chain mechanism,
–r
O3
= k
1
10
pH 14
C
O3
+ k
2
10
1/2(pH 14)
C
O3
1.5
; data
on k determined
Reaction H
2
O
2
–O
3
also studied
19
(1935)
Batch reactor, spectrophotometric, and

iodometric analysis
0 and 27ºC, pH 0.7–2.8, HClO
4
Chain mechanism
–r
O3
= k10
1/2(pH 14)
C
O3
49
(1950)
Batch reactor, pH 7.6–10.4, T: 1.2–19.8ºC –r
O3
= kC
O3
; data on k determined 50
(1954)
Batch reactor, cell, saturation reactor,
Spectrophotometric, iodometric, manometric
analysis, 25ºC, pH 0.2–10, HClO
4

(10
–5
–0.96 M), H
2
SO
4
, NaOH, Cu(II), Fe(II)

of perchlorate, glicine, phosphate, arseniate
buffers
Influence of different acids and salts
Kinetics varies depending on pH values
For acid–neutral pH: –r
O3
= kC
O3
3/2
For pH > 7: –r
O3
= kC
O3
2
Henry constant also determined
51
(1956)
Semibatch bubble column. Iodometric analysis.
10–40ºC, pH 2–4 (H
2
SO
4
), pH 6 (phosphate
buffer), pH 8 (borate buffer)
Kinetics depends on pH:
pH 2–4: –r
O3
= kC
O3
2

, pH 6: –r
O3
= kC
O3
2–1.5
pH 8: –r
O3
= kC
O3
Ozone–secondary wastewater also studied
52
(1971)
Semibatch bubble column until saturation.
pH: 0.22–1.9, T: 5–40ºC
–r
O3
= kC
O3
or –r
O3
= kC
O3
2
. Chain
mechanism. Data on k determined
53
(1971)
Batch reactors, spectrophotometric analysis
25ºC, pH 2–11 Different organics and salts
Influence of organics

Relative rates of organic decomposition
Two ways of ozone decomposition, critical
pH
54
(1976)
Semibatch packed column (95 cm, 3.7 cm
diam). Packed: glass rings. Iodometric
analysis, 27ºC, pH 8.5–13.5, KOH,
k
L
a and a determined from CO
2
absorption in
arseniate solutions
–r
O3
= k10
pH 14
C
O3
,
55
(1976)
3.2 L Batch reactor, Iodometric analysis,
3.5–60ºC, pH 0.45–12, phosphate buffers
–r
O3
= k10
0.123(pH 14)
C

O3
56
(1979)
Cell reactors, spectrophotometric analysis,
25ºC, pH 1–3 (H
2
SO
4
), 4–5 (acetate buf),
7–9.5 (phosphate-borate buf), 10–11.5
(carbonate buf.), 12–13.5 NaOH
For pH < 8 no sufficient reliable data was
obtained to fix the ozone reaction order
For pH > 8, –r
O3
= k10
0.88(pH 14)
C
O3
Reactions O
3
-CN
–
and O
3
dyes were also
studied
57
(1981)
Cell reactor. Spectrophotometric analysis. 20ºC,

pH 3, 7, and 9, H
2
SO
4
, and NaOH
Kinetics depends on pH:
pH 3: –r
O3
= kC
O3
3/2
, pH 7: –r
O3
= kC
O3
3/2 to 1
pH 9: –r
O3
= kC
O3
Kinetics of O
3
-bromide studied
58
(1981)
©2004 CRC Press LLC
TABLE 2.7 (continued)
Works on Aqueous Ozone Decomposition Kinetics
Reacting System and Operating Conditions Main Observations Reference #
2.8 L Continuous stirred batch reactor,

spectrophotometric analysis, 20ºC, tap water
(pH not given)
Mass transfer and chemical reaction model
proposed and tested.
Observed zero order kinetics for ozone
Mass transfer coefficient also determined
59
(1982)
Spectrophotometer cell, spectrophotometric
analysis. 20ºC, pH 11–13, NaOH, carbonate
and acetate (in some cases)
Acetate and carbonate inhibitors
Hydrogen peroxide promoter
Mechanism, –r
O3
= k10
(pH 14)
C
O3
O
3
–
• and CO
3
–
• identified
60
(1982)
Cell reactor, spectrophotometric analysis, 20ºC,
pH 8–10

Phosphate buffer, carbonate (in some cases)
Carbonate and phosphate inhibitors,
Hydrogen peroxide also investigated.
Mechanism and kinetics: –r
O3
= k10
(pH 14)
C
O3
20
(1982)
Batch reactor for k determination. 20ºC, pH < 4
H
2
SO
4
, pH 6–9 phosphate buf., pH 7–10
borate buf., Colorimetric analysis. Semibatch
bubble column for checking results
Dependency on pH:
pH < 4, –r
O3
= kC
O3
2
, pH > 4, –r
O3
=
k10
0.55(pH 14)

C
O3
2
Kinetics checked from O
3
absorption in water
in bubble column. Henry constant and mass
transfer coefficient also determined
61
(1982)
Comparison of results from previous works
3.5–60ºC, pH 0.45–10.2
Different ozone reaction orders were tested.
No statiscal differences between ozone
reaction orders (1st, 2nd and 3/2). Data can
be fitted with an ozone first-order kinetics.
Potential effects of impurities and buffer
systems used
62
(1983)
Pulse radiolysis experiments, 21ºC, pH 6.3–7.9,
phosphate buffers adjusted with NaOH and
HClO
4
Studies of free radical formation and decays
(O
3
–
•, HO
3

•). Initiation and propagation
steps of ozone decomposition. Influence of
phosphate
22
(1984)
2nd part of Reference 22 Studies of termination chain reactions, The
role of HO• and HO
4
• radicals. Full
mechanism proposed. Recombination of
2HO
4
• most probable termination reaction
23
(1984)
Batch reactors, conditions as in Reference 52.
Spectrophotometric and colorimetric analysis
Definition of promoters, initiators and
inhibitors of ozone decomposition in water.
First-order reaction. Full mechanism
proposed (see Table 2.3). Full kinetic
expression see Equation (2.68)
(21)
1985
Stopped flow cell. Spectrophotometric
analysys. 20ºC, pH 12 NaOH, Na
2
CO
3
effect

Kinetics at alkaline conditions,
pH 12: –r
O3
= k
1
C
O3
+ k
2
C
O3
2
, pH>8: –r
O3
=
kC
O3
in the presence of carbonates. Full
mechanism proposed. Formation of HO
2
–

and O
2
in the initiation reaction
24
(1985)
©2004 CRC Press LLC
TABLE 2.7 (continued)
Works on Aqueous Ozone Decomposition Kinetics

Reacting System and Operating Conditions Main Observations Reference #
750 ml Batch stirred reactor,
Spectrophotometric analysis. 10–40ºC,
pH 2.5–9, phosphate buffers
Mechanism proposed and tested.
Kinetics: –r
O3
= k
1
C
O3
+ k
2
10
1/2 (pH 14)
C
O3
1.5
Determination of rate constants and energy of
activation
46
(1987)
Pulse radiolysis experiments. 20ºC, pH 4.5–9
Acetate buffers
Acetate acts as inhibitor of ozone
decomposition
Identification of acetate, ozonide free
radicals
Mechanism proposed. Determination of rate
constants

63
(1987)
Treatment of data from previous works Discussion on the chain termination reaction
Termination proposed: O
3
–
•+O
2
–
•
Kinetics: 3 < pH < 7: –r
O3
= k
1
10
1/2(pH 14)
C
O3
3/2
,
pH = 7: –r
O3
= k
1
10
(pH-14)
C
O3
3/2
, 7 < pH < 12:

–r
O3
= k
1
10
1/2(pH-14)
C
O3

64
(1988)
Batch stirred reactor. Colorimetric analysis.
10–30ºC, pH 2–8.5: phosphate buf., pH 6:
Carbonates, chlorides, sulphates and salt-free
water
Full mechanism proposed. Dependency on
pH, nature of salts, and ionic strength. Ozone
reaction orders vary from 1st, half, three
halves, and 2nd.
Rate constants are determined from integral
methods.
65
(1989)
2-cm path quartz cell reactor.
Spectrophotometric analysis. pH > 12, 20ºC,
NaOH. Carbonates: 0–0.05M
Stability of ozone and no effect of carbonates
at pH > 12. Mechanism proposed and a new
step : reaction HO• + OH
–®

O
–
+ H
2
O. Free
radicals such as O
2
–
• and O
3
–
•. Process
simulation.
66
(1989)
100 ml glass syringes. Spectrophotometric
analysis. Oxygen measurements by GC.
0–46ºC, pH 0–4, HClO
4
Formation of oxygen is proposed and
confirmed by isotopic interchange. Primary
step: dissociation of oxygen at acid pH.
Ozone first-order kinetics. Activation
energy: 79.5 kJmol
–1

67
(1991)
Batch–semibatch photoreactors. 5.3 W low
pressure Hg lamp. Colorimetric analysis,

20ºC, pH 2–10, NaOH, and H
2
SO
4
continuous
feeding.
Kinetics: No UV: –r
O3
= k
1
10
0.395(pH 14)
C
O3
1.5
With.UV:–r
O3
= k
2
10
0.064(pH 14)
C
O3
1.5
I
0.9
+
k
1
10

0.395(pH 14)
C
O3
1.5
Chloride and nitrate: weak scavengers of HO•
68
(1996)
Pulse radiolysis experiments. 20ºC, pH > 11 Kinetic modeling with ozone decomposition
mechanisms of Reference 21 and Reference
24 with some minor modifications.
Inorganic carbon reactions included.
Mechanism of Reference 24 better at high
pH
69
(1992)
©2004 CRC Press LLC
References
1. Riebel, A.H. et al., Ozonation of carbon-nitrogen bonds. I. Nucleophilic attack of
ozone, J. Am. Chem. Soc., 82, 1801–1807, 1960.
2. Buxton, G.V. et al., Critical review of data constants for reactions of hydrated elec-
trons, hydrogen atoms, and hydroxyl radicals (.OH/.O
-
) in aqueous solution, J. Phys.
Chem. Ref. Data, 17, 513–886, 1988.
3. Ayres, G.H., Análisis Químico Cuantitativo, Ediciones del Castillo, Madrid, 1970.
TABLE 2.7 (continued)
Works on Aqueous Ozone Decomposition Kinetics
Reacting System and Operating Conditions Main Observations Reference #
Batch reactor. Natural organic matter (NOM):
0–0.25 mM as organic carbon.

Spectrophotometric analysis with NOM-free
water, Colorimetric analysis: with NOM in
water. pH 7.5 phosphate buffer, t-butanol, and
p-chlorobenzoate in some runs
Kinetic modeling of ozone decomposition
following mechanisms in Reference 21 and
Reference 24. Predictions of free radical and
molecular ozone concentrations with time.
Proposition of some HO+NOM reactions to
fit experimental results
25
(1997)
100 ml glass syringe reactors. 22–45ºC, pH 0–4,
HClO
4
. Presence of H
2
O
2
produced from γ-ray
irradiation of oxygen. Spectrophotometric
analysis
Mechanism proposed and tested.
Formation of H
2
O
2
checked. Initiation and
termination reactions are surface-catalyzed
70

(1998)
Semibatch bubble photoreactor.
Spectrophotometric analysis. Low pressure Hg
lamp. 20ºC, pH 7
Kinetic modeling of ozone decomposition by
UV radiation. Determination of the ozone
quantum yield: 0.64
71
(1996)
1 cm stirred spectrophotometric quartz-glass
cells. Spectrophotometric analysis. Batch
flask, Colorimetric analysis. Activated carbon
or black carbon, Different organics: methanol,
acetate, p-chlorobenzoate, NOM, etc.
Catalytic effect of activated carbon and a
carbon black. Effects of promoters and
inhibitors of ozone decomposition. Ozone
first-order kinetics.
Catalytic reaction generates hydroxyl radicals
27
(1998)
500 ml graduated glass cylinder. 20ºC, pH 7.5
phosphate buffer. NOM from 18 surface
waters. p-chlorobenzoate addition to
determine HO concentration
27

Promoting and inhibiting effects of NOM.
Ozone decomposition in two steps: t < 1 min
and t > 1 min: ozone first-order kinetics. Rate

constant correlated to physical parameters of
surface waters
44
(1999)
Stopped flow spectrophotometer.
Spectrophotometric analysis. 25ºC,
pH 10.4–13.2 NaOH, 0.5 M ionic strength
with HClO
4
Use of mechanism in Reference 24 with some
modifications. Small amounts of H
2
O
2

adsorbed on the spectrophotometer cell can
alter the kinetic results. Prediction of ozone
and other free radicals half live
72
(2000)
38.3L-batch reactor, Liquid ozone UV
photometry analyzer, 18–30ºC, 2.6 and 7,
phosphate buffer. Ionic strength: 0.01 M
Mechanism and rate equation as in
Reference 45
Rate constant and activation energies
determined
73
(2002)
©2004 CRC Press LLC

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