C Max. O3-charged

wt% Vol. % C at STa

[g Nm–3]

C at I NDb

[g m–3]

For 1 kg

[O3 odt–1]

[kg O3 odt–1]

5.0 3.4 72.6 491 0.17 1.4

6.0 4.1 87.4 591 0.15 1.7

6.8 4.7 100.0 676 0.13 1.9

7.0 4.8 102.3 692 0.13 1.9

8.0 5.5 117.3 793 0.12 2.2

9.0 6.2 132.5 896 0.10 2.4

10.0 6.9 147.7 999 0.09 2.7

10.2 7.0 150.0 1014 0.09 2.7

11.0 7.6 163.0 1102 0.09 2.9

12.0 8.3 178.5 1207 0.08 3.2

13.0 9.1 194.0 1312 0.07 3.4

13.4 9.3 200.0 1352 0.07 3.5

14.0 9.8 209.7 1418 0.07 3.7

15.0 10.5 225.4 1524 0.06 3.9

16.0 11.3 241.3 1632 0.06 4.2

16.5 11.7 250.0 1691 0.06 4.3

19.6 14.0 300.0 2029 0.05 5.1

a. ST = standard conditions: T0 = 273.15 K, P0 = 101.3 kPa.

b. IND = industrial conditions: T = 323.15 K, P = 810.6 kPa = 8 bar.

c. X g = V g/(V g+ V L) at 10% pulp consistency and IND.

d. Assuming an upper limit of X g = 0.25 to obtain a reasonably

high ozone consumption rate.

7.5.4

Chemistry of Ozone Treatment

Manfred Schwanninger

Ozone treatment is a very effective way to remove residual lignin that remains

after pulping. The structure and reactivity of the residual lignin have already been

described (see Section 7.3.2.2). One of the major disadvantages of ozone as a

7.5 Ozone Delignification 785

bleaching agent is its moderate stability in aqueous solutions [1–4]. It has the tendency

to decompose in water, generating some very reactive, highly unselective

radical species [1–5]. Hydroxide ions are known to catalyze ozone decomposition

and to promote hydroxyl radical formation [1,3,4,6], whilst another drawback is

the undesired degradation of cellulose [7–23].

Ozone Decomposition

The pathways and kinetics of the decomposition of aqueous ozone are of interest

for a wide range of topics, not only for pulp bleaching, and have therefore been

studied intensively [1–5,24]. The chain mechanism of ozone decomposition

(Scheme 7.28) is based on the studies of Buhler et al. [5], while Staehelin et al. [1]

showed the decomposition of ozone and the formed intermediates (Scheme 7.28).

O2

-

HO2

O2

HO4

O3

OH

HO3

O3

-

1O2

O3

H2O

OH-

O2

H2O

H+

HO4

+

H2O2 + 2O3

HO3 H2O2 + O3 + O2

+

Termination

HO4

HO4

Scheme 7.28 Chain mechanism of ozone decomposition

according to Staehelin et al. [1].

The decomposition occurs by a radical chain mechanism, which in pure water

is initiated by the reaction between hydroxide ions (OH–) and ozone [Eq. (89). Superoxide

_O2

– then reacts with ozone rather selectively as part of a chain cycle [1].

O3_OH_→HO2 _ _ _O_2 _89_

The hydroxide-ion-catalyzed decay of ozone is expressed in the following general

rate equation (1):

d _ O 3

dt _ _ k _ _ O 3 a _ _ OH _ b _90_

The reaction order b with respect to hydroxide ion concentration varies from

0.36–1.0 and the reaction order a for ozone is reported to vary between 1.5–2.0 [aa,

bb]. Pan et al. have studied the decomposition rate of ozone in a pure aqueous

786 7Pulp Bleaching

7.5 Ozone Delignification 787

2 4 6 8

second order reaction rate

O

-concentration [mg/l]

Time [min]

Second Order Rate Constant

[mol-1*l*min-1]

pH

0 5 10 15 20 25 30

[O

] at pH 3 [O

] at pH 7

Fig. 7.82 Effect of pH on second order rate constant of ozone

decomposition and the course of ozone concentration in aqueous

solution at 25 °C according to Pan et al. [25].

system as a function of pH at 25 °C. The results revealed a rapid increase of the

second order rate constant from pH 4 to 7 as seen in Fig 7.82.

The concentration of dissolved ozone decreases rapidly at a neutral pH, while it

remains quite stable under acidic conditions (pH 3) within a time frame relevant

for industrial ozone bleaching applications.

The first propagation step of the chain reaction proposed by Weiss for the

decomposition of aqueous ozone can be described with the intermediates _O3

–

and HO3_ [5]. The elementary reactions of these species and their constants are

presented in Eq. (91) [5]:

O3__O_2 →_O_3 _1O2 k _ 1_6 _ 109M_1s_1

_O_3 _H_

ka

kb

HO3 _ ka _ 5_2 _ 1010M_1s_1 kb _ 3_7 _ 104M_1s_1

HO3 _→HO_ _ O2 k _ 1_1 _ 105M_1s1_1

_91_

In water, the decay of HO3_ is very rapid, with a half-life of about 6 ls. Since

_O2

– reacts rapidly with ozone, but relatively slowly with organic compounds, the

latter will interfere with ozone decomposition in aqueous solutions by scavenging

OH radicals rather than _O2

– [5]. The second propagation step and the reactions of

these species and their constants are presented in Eq. (92):

HO_ _ O3→HO4 _ k _ 2_0 _ 109M_1s_1

HO4 _→HO2 _ _ O2 k _ 2_8 _ 104M_1s_1

HO2 _ _ _O_2 _H_ p K a_ 4_8

_92_

The transient HO4_ might be a charge-transfer complex (HO_O3) [1]. The lifetime

of HO4_ was found to be much longer than its accumulation rate, and therefore

it acts as a carrier reservoir within the chain cycle. As a consequence, HO4_

is the important transient for chain termination reactions (Fig. 7.82). The termination

reactions shown in Eq. (93) are the dominating ones in the presence of

high ozone concentrations [1].

HO4 _ _ HO4 _→H2O2_2O3

HO4 _ _ HO3 _→H2O2_O3_O2 _93_

If, at low ozone concentrations, organic solutes are present, their dominant

effect will be to withdraw OH radicals; at high ozone concentrations this will similarly

occur with HO4_. Some organic materials are thereby able to regenerate _O2

–

in order to sustain the chain.

As the initial step of the ozone decomposition is the reaction between ozone

and the hydroxide anion [Eq. (89) and Scheme 7.28], strong pH dependence is

expected [6]. Gurol and Singer [4] investigated the kinetics of ozone decomposition

in the pH range from 2 to 10. They found that ozone decomposes rapidly at a pH

above about 6.5, but remains quite stable under acidic conditions; indeed, this

finding has been confirmed by several authors [6,25,26]. Hydrogen peroxide, also

used as a bleaching chemical, is incapable of initiating ozone decomposition;

however, its deprotonated form, the hydroperoxy anion (HO2

–) has such ability [6].

At pH < 12 and hydrogen peroxide concentrations >10–7 mol L–1, HO2

– has a

greater effect on the decomposition rate than the hydroxide anion (OH–) [3], and

this might be important for bleaching sequences. In the radical-type chain reaction

decomposition of ozone, inorganic and organic compounds can be divided

into three categories: (a) initiators; (b) promoters; and (c) inhibitors [6,24,27,28].

Initiators are substances that are capable of initiating the decomposition of ozone

to the superoxide anion radical [Eq. (89)], while promoters are radical converters

forming the superoxide anion radical from the hydroxyl radical (Scheme 7.29).

Inhibitors are substances that react with the hydroxyl radical without the formation

of superoxide anion radical, called radical scavengers, such as bicarbonate

and carbonate, leading to the corresponding radicals [24]. Some examples are given

in Tab. 7.37.

CH3OH

OH H2O

H2C

OH

H2C

OH

O2

OO

O2

-

B- BH

H2CO

Scheme 7.29 Reaction of methanol as a typical hydroxyl radical

to superoxide anion radical converter acting as a promoter [24].

788 7Pulp Bleaching

Tab. 7.37 Typical initiators, promoters, and inhibitors for

decomposition of ozone by radical-type chain reactions

[6, 24–28].

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