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ABSTRACT

Ozone is a powerful oxidant used in the water-treatment industry to make drinking water wholesome and to remove undesirable taste, odor and color. In order to evaluate such treatment processes, it is imperative to understand the decomposition

kinetics of aqueous ozone. The decomposition of aqueous ozone occurs via chain reactions involving hydroxyl radicals. Many dissolved substances in water react with the hydroxyl radicals and alter the rate of aqueous ozone decomposition. This study examines the effects of pH, carbonate and humic substances on aqueous ozone decomposition. Batch kinetic experiments were conducted at 22 "C in solutions containing 0.01 M total phosphate species and total

carbonate species concentration varying from 0 to 10 x 10'' M. The

experimental data suggest that ozone decomposition in the presence of carbonate species conforms to simple first-order kinetics. The rate of ozone decomposition increased with increase in pH and was retarded with increase in carbonate species concentration, due to scavenging of the hydroxyl radicals. Ozone decomposition in the

presence of humic substances was characterized by an initial immediate rapid ozone utilization followed by first-order depletion

of ozone. Similar behavior was observed in natural waters. The

pseudo-first-order rate constants decreased with increase in

initial ozone dose and decreased with increase in initial concentration of humic material. The Staehelin and Hoigne (1985)

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involved during ozone depletion. A correlation reported in the

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ACKNOWLEDGEMENTS

This project was funded by the American Water Works

Association Research Foundation.

I would like to thank my advisor, Dr. Philip Singer for

the contribution of his patient support and expertise. I also owe thanks to Mr. Chris Hull for all his ideas and assistance.

Additional thanks to Dr. Mark Sobsey and his group of students for their help in the form of computer time and facilities.

Mr. Richard Hall's assistance with the fabrication of the

experimental set-up is greatly appreciated. I also extend my thanks to Mr. Tony Greiner for helping me with the analysis of

natural water samples.

Most of all, I thank my parents, for their encouragement

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TABLE OF CONTENTS

Page

List of Tables... vi

List of Figures... ix

1. Introduction ... 1

2. Theoretical Background ... 5

A. Aqueous Ozone Decomposition ... 5

(i) Effect of pH and Temperature... 5

(ii) Effect of Humic Material... 13

B. Kinetic Models for Ozone Decomposition .... 17

(i) Ozone Decomposition in Pure Water ... 17

(ii) Aqueous Ozone Decomposition in the Presence of Humic Material ... 19

3. Experimental Methods ... 23

A. Introduction... 23

B. Experimental Set-up... 23

C. Experimental Procedure ... 27

(i) Initial Preparation ... 27

(ii) Test Solution Preparation and Sampling Methods ... 29

D. Analytical Procedures for Aqueous Ozone Measurement... 32

(i) lodometric Method... 32

(ii) UV Method... 33

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Page

4. Results and Discussion... 41

A. Introduction... 41

B. Effect of pH ... . ... .... 41

C. Effect of Carbonate and Bicarbonate Ions ... 56

D. Effect of Hvunic Material... 72

E. Ozone Depletion in Natural Waters ... 94

F. Engineering Applications of Results ... 102

5. Conclusions... 104

6. References... 107

7. Appendices ... 112

Appendix A Effect of Quality of Buffer

Chemicals on Ozone Decomposition 112

Appendix B Derivation of the Mixed-Order Rate

Equation Using the Staehelin and Hoigne

Mechanism 114

Appendix C Derivation of the Generalized Mixed-Order

Rate Equation Using the Staehelin and

Hoigne Mechanism 117 Appendix D Data Analysis Procedures and Sample

Calculations 120 Appendix E

Appendix F

Derivation of the Rate Equation for Ozone

Depletion in the Presence of Humic

Substances Using the Staehelin and Hoigne

Mechanism 128

Results of Selected Experiments on Dis¬

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Table 2. 1

Table 2. 2

Table 2 3 Table 2 .4

Table 3 .1

Table 4 .1

Table 4 .2

Table 4 .3

Table 4.4

Table 4 .5

Table 4 .6

Table 4 .7

Table 4 .8

Table 4 9

Table 4. 1

List of Tables

Summary of aqueous ozone decomposition kinetics Summary of rate equations for dissolved ozone decomposition

Fundamental reactions of ozone decomposition (after Staehelin and Hoigne (1982, 1985)) Fundamental reactions of ozone decomposition

(after Tomiyasu, Fukutomi and Gordon (1985))

Experimental Matrix

Comparison of theoretical and experimental k,' in

mixed-order model for experiments at 22'C with no inorganic & organic carbon

Page 6 8 11 20 24 47 Summary of rate constants for first and second-order models for experiments at 22"C with no inorganic

and organic carbon 47

Comparison of theoretical and experimental k,' in

mixed-order model for experiments at 22°C with no

organic carbon 58

Summary of rate constants for the first-order model for experiments at 22"C with no organic carbon 62 Summary of rate constants for the second-order model for experiments at 22'C with no organic carbon 64

Kinetic parameters for ozone decomposition in the

presence of dissolved humic material 77

Estimated (k'p + k',) and k',k'p from measured

pseudo-first-order rate constant (K,') of

humic material 77

Estimated k'p from measured pseudo-first-order rate

constant (K*^) for ozone decay in the presence of

humic material 85 Comparison of estimated and measured pseudo-first

order rate constant (K',) for ozone decay in the

presence of humic material 85 10 Effect of initial ozone dose on the actual amount

ozone consumed during ozone decay in the presence

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Page

Table 4.11 Effect of initial ozone dose on the kinetic

parameters for ozone decomposition in the presence

of dissolved humic material 88 Table 4.12 Effect of initial ozone dose on the kinetic

parameters for ozone consumption in natural waters 101

Table 4.13 Comparison of estimated and measured pseudo-first

order rate constant (K'^) for ozone consumption in

natural waters 101

Table B.l Fundamental reactions of dissolved ozone

decomposition (after Staehelin and Hoigne

(1982,1985)) 114

Table C.l Reactions of dissolved substances with ozone and

hydroxide radical formed during ozone decomposition 117

Table D.l Experimental k'g values for experiments conducted

at 22'C, pH 5.5 with 0.01 M total phosphate and no

inorganic & organic carbon 123

Table D.2 Equilibrium constant correction factor for ionic

strength effects. (Using Equation (B.IO)) 123

Table D.3 Theoretical k'g values for experiments conducted

at 22''C, pH 5.5 with 0.01 M total phosphate and no

organic carbon 123

Table D.4 Experimental k'j values for experiments conducted

at 22°C, pH 5.5 with 0.01 M total phosphate and

0.002 M total carbonate and no organic carbon 126

Table D.5 Dissociation fractions of phosphoric acid at

relevant pH values at 22 "C with /i = 0.01 127

Table D.6 Dissociation fractions of carbonic acid at

relevant pH values at 22 "C with /x = 0.01

in water 127 Table E.l Reactions of dissolved humic substances with ozone

and hydroxide radical formed during ozone decomposi¬

tion in water 128 Table F.l Residual ozone concentration versus time data

illustrating the effect of pH on ozone decomposi¬

tion experiments conducted at 22°C with no

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Page Table F.2 Residual ozone concentration versus time data

illustrating the effect of pH on ozone decomposi¬

tion in experiments conducted at 22"C with

0.002 M C^ and no organic carbon 134

Table F.3 Residual ozone concentration versus time data

illustrating the effect of C^ on ozone decomposi¬

tion in experiments conducted at 22'C and pH 7.0

with no organic carbon 134 Table F.4 Residual ozone concentration versus time data

illustrating the effect of Cj, on ozone decomposi¬

tion in experiments conducted at 22°C and pH 7.0

with no organic carbon 135

Table F.5 Residual ozone concentration versus time data illustrating the effect of humic acid content on ozone depletion in experiments conducted at 22°C

and pH 5.5 with 0.002 M C^ ' 135

illustrating the effect of humic acid content on ozone depletion in experiments conducted at 22'C

and pH 7.0 with 0.002 M C^ 136

illustrating the effect of C^ on ozone depletion

in experiments conducted at 22"C and pH 7.0 with

0.002 M C^ 137 Table F.8 Residual ozone concentration versus time data

illustrating the effect of C^ on ozone depletion

in EBMUD water at 18"C and pH 7.6 with 3.90 mg/L

TOC and alkalinity 60 mg/L 137

illustrating the effect of C^ on ozone depletion

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Figure 2.1 Figure 2.2 Figure 3.1 Figure 3.2 Figure 3.3 Figure 3.4 Figure 3.5 Figure 4.1 Figure 4.2 Figure 4.3 Figure 4.4 Figure 4.5 Figure 4.6 Figure 4.7 Figure 4.8 Figure 4.9

List of Figures

Reactions of dissolved ozone in pure water (after Staehelin and Hoigne, 1982)

Reactions of dissolved ozone with humic material

(after Feng and Legube, 1991) Schematic of Experimental Set-up

Schematic of Experimental Batch Reactor Calibration Curve I of indigo method for 0-6 mg/L dissolved ozone

Calibration Curve II of indigo method for

0-3 mg/L dissolved ozone

Calibration Curve III of indigo method for 0-1 mg/L dissolved ozone

Page 10 14 26 28 38 39 40 Decomposition of ozone in water at pH 2.0,

demonstrating no ozone loss due to volatilization 42

Effect of pH on ozone decomposition in water

(first-order plots) 50

Effect of pH on first order rate constant of ozone

decomposition in water 52

(second-order plots) 53 Comparison of first and second-order models to

describe ozone decomposition in water 54

Effect of initial ozone dose on the rate of ozone

decomposition in water at pH 7.0 55

containing 0.002 M total inorganic carbon

(first-order plots) 59 Effect of pH on ozone decomposition in water

containing 0.002 M total inorganic carbon (second-order plots) 60 Comparison of first and second-order models to

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Page Figure 4.10 Effect of initial ozone concentration on the rate

of ozone decomposition in water containing

inorganic carbon at pH 7.0 66

Figure 4.11 Effect of total inorganic carbon content on ozone

Figure 4.12 Effect of total inorganic carbon content on ozone

Figure 4.13 Effect of inorganic carbon content on first order

rate constant for ozone decomposition in water

with varying pH 71 Figure 4.14 Effect of total organic carbon content on ozone

decomposition in water pH 5.5 73 Figure 4.15 Effect of total organic carbon content on ozone

decomposition in water at pH 7.0 76 Figure 4.16 Effect of total organic carbon content on initial

ozone consumption in water at pH 7.0 79 Figure 4.17 Effect of initial ozone dose on ozone decomposition

in water with an organic carbon content of 2 mg/L 87

Figure 4.18 Effect of initial ozone dose on ozone decomposition

in water with an organic carbon content of 4 mg/L 90 Figure 4.19 Effect of initial ozone dose on initial ozone

consumption in water with an organic carbon

content of 4 mg/L 91 Figure 4.21 Effect of initial organic carbon content on

long-term ozone consumption at pH 7.0 93

Figure 4.21 Effect of ozone dose on ozone consumption at

pH 7.0 95

Figure 4.22 Effect of initial ozone dose on ozone constimption in a natural water (EBMUD) of pH 7.3 97 Figure 4.23 Effect of initial ozone dose on ozone consumption

in a natural water (EBMUD) of pH 7.7 98 Figure 4.24 Effect of initial ozone dose on ozone consumption

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Page

Figure A.l Effect of quality of buffer chemicals on ozone

Figure D.l Variation of experimenatal kg' with time

in experiments conducted at pH 8.5 122 Figure E.l Reactions of dissolved ozone in the presence of

solute HA which react with ozone or which interact

with hydroxyl (OH) radicals by scavenging and/or converting OH into HOg (after Staehelin and

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1. INTRODUCTION

The main purpose of drinking water treatment is to make

the water safe to drink. Outbreaks of waterborne diseases such as

typhoid, cholera, hepatitis etc. were widespread until water

treatment came into being approximately a century ago. Chlorine

has been the most common disinfectant used to quell the outbreaks

of such waterborne diseases.

Over the last two decades, there has been growing concern

about the interactions between chlorine and naturally-occurring

organic matter in raw water to produce by-products which are

suspected carcinogens. This grave concern has caused the water

industry to examine and use other disinfectants, including ozone. Ozone, apart from being a powerful drinking water disinfectant, has been found to oxidize (i) iron and manganese,

(ii) taste and odor-causing compounds, and (iii) precursors of

chlorination by-products, and (iv) to enhance particle removal.

These advantages have made a number of water treatment plants in North America opt to use ozone.

In 1974, the Safe Drinking Water Act (SDWA) was passed to

ensure that all drinking water supplied to the general public is

safe and meets certain standards of quality. The SDWA was amended

and new guidelines were set in 1986. Among the new guidelines, the

CT product is a parameter used to demonstrate that a particular

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the Surface Water Treatment Rule (SWTR) lists the CT values for

these microorganisms. CT is the product of the concentration of

the disinfectant in mg/L and the time, in minutes, that the

disinfectant is in contact with the water.

In order to evaluate the CT values in full-scale ozone

contactors, it is imperative to mathematically model the residence time distribution and concentration profiles in the ozone contacto¬ rs. Such a model would enable easy prediction of the residual ozone concentrations and the corresponding residence time, thus making the calculation of the CT values simple and inexpensive. Using the predicted CT values, concentration and residence time distribution, it would be possible to design optimal ozone

contactor configurations.

The residual liquid-phase ozone concentration at any time depends significantly on the rate of ozone transfer from gas to liquid phase and the rate of ozone depletion in water. In order to estimate the rate of ozone depletion in water, batch ozone kinetic

experiments need to be conducted. Using the results of batch ozone decomposition studies and the model of the ozone contactor, it

should be possible to determine the ozone dose required to achieve a target ozone residual concentration in a given water or to comply

with ozone disinfection regulatory requirements.

Ozone is unstable in water. Ozone decomposes via complex

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react directly with dissolved substances in water. Depending on

the water composition, the hydroxyl and superoxide radicals may

themselves react with solutes leading to different pathways hence,

different types of kinetics. Several researchers have studied the

decomposition kinetics of aqueous ozone in batch reactors and it is

known from these studies that the rate of ozone decomposition

increases with increase in pH. The presence of radical scavengers

such as carbonate and bicarbonate ions lowers the rate of ozone

decomposition and enhances its stability. Organic impurities such

as aquatic humic substances act both as initiators and promoters of

the chain reaction and as radical scavengers.

The objectives of this study were to: (i) quantify the

impact of pH, carbonate content (alkalinity), humic substances and

initial ozone dose on the kinetics of ozone decomposition, (ii)

using results of the above objective, evaluate mechanistic models

of ozone decomposition developed in the literature and estimate the

kinetic parameters of such models, and (iii) utilize the above

results to compare and model the rate of ozone depletion in natural

waters. It is important to reiterate that, though many earlier

studies have tried to address parts of the above objectives, this

investigation was uniquely designed to allow for integration of the

results of all the objectives into a mathematical model describing

the residence time distribution and ozone concentration profiles in

ozone contactors. As mentioned earlier, the mathematical model

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optimization of initial ozone dosages and flow configurations in

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2. THEORETICAL BACKGROUKD

In order to predict the effectiveness of ozone as an oxidant and disinfectant in both drinking water and wastewater treatment systems, a rate expression describing the aqueous decomposition of ozone is needed. Such an expression would also be helpful in designing and evaluating the aforementioned treatment systems. A number of investigators have developed such rate expressions. This

section reviews these investigations. A. Ac[ueous ozone decomposition

(i) Effect of pH and Temperature

The decomposition rate of dissolved ozone is consider¬

ably affected by temperature and pH. Several researchers have studied the effect of these variables. Table 2.1 summarizes the

results of these studies. It shows that there is disagreement among researchers with regard to the order of the reaction and the magnitude of the rate constant. However, all investigators agree that decomposition of ozone is catalyzed by the hydroxide ion and that the reaction rate drops almost to zero below pH 2. Table 2.2 reports the rate equations obtained in some of these studies. These equations clearly show the pH (or hydroxide ion concentra¬ tion) dependence of the ozone decomposition rate.

In order to comprehend the effect of pH on ozone

decomposition, it is important to understand the underlying

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Table 2.1 Summary of aqueous ozone decomposition kinetics

Reference Date _PH

Reaction order

Temp. with respect

("C) to ozone

Rothmund,et al.^ (1913) 2 - 4 0 2

Sennewald' ₍₁₉₃₃₎ 5.8 - 8 0 2

Weiss^ ₍₁₉₃₅₎ 2 - 8 0 _3/2

Alder & Hill^ ₍₁₉₅₀₎ 1 - 2. 8 0.0 ͣ - 27 1

Stumm' ₍₁₉₅₄₎ 7.6 - 10. 4 1.2 •-19.8 1

Kilpatrick^ (1956) 0.0 - 6. 8 25 _3/2

Kilpatrick^ (1956) 8.0 - 10 25 2

Rankas, et al.^ (1962) 5.4 - 8. 5

-25 _3/2

Czapski^ (1968) 10 - 13 25 1

Rogozhkin^ (1970) 9.6 - 11.9 25 1

Hewes & Davidson^ ₍₁₉₇₁₎ 6 10 - 50 _3/2-2

Hewes & Davidson^ ₍₁₉₇₁₎ 8 10 - 20 1 Hewes & Davidson^ ₍₁₉₇₁₎ 2 - 4 30 - 60 2 Merkulova, et al.^ (1971) 0 - 2 5 - 40 1-2

Shambaugh & Melynk^ (1976) 9 20 1

Rizutti, et al.^ (1976) 8.5 - 13. 5 18 - 27 1

Kuo, et al. (1977) 2.7 - 11 25 _3/2

Chang^ (1978) 9 23.5 1

Morooka, et al.^ (1978) 2 - 9 3 - 35 _3/2-1

Sullivan & Roth' ₍₁₉₇₉₎ 0.5 - 10 3.5 - 60 1 Teramoto, et al.^ (1981) 1 - 13. 5 25 1

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Table 2.1 continued

Reference Date _PH

Forni, et al.-* (1982) 12 Staehelin & Hoigne (1982) 8-10 Sheffer & Esteirson' (1982) 6-8

Tomiyasu, et al. (1985) 12

Sotelo, et al. (1987) 2.5-9

Yurteri & Gurol (1988) 6.8 - 0

Grasso & Weber (1989) 5-9

Reaction order

Temp. with respect

('C) to ozone

20 1

19 - 22. 5 0

20 1-2

10 - 40 _1-3/2 20 ± 1 1-2

Not reported 1-2

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Table 2.2 Summary of rate equations for dissolved ozone decompo¬

sition

Reference Date _{Rate equation r = -——z—} at

Sennewald'

Weiss^

Alder & Hill^

Stumm'

Rizutti, et alJ (Packed Column)

Morooka, et al.^

Sullivan & Roth^

Gurol & Singer

Staehelin & Hoigne

Forni, et al.'

Tomiyasu, et al.

Sotelo, et al. Yurteri & Gurol Grasso & Weber /

(1933) (1935) (1950) (1954) (1976) (1978) (1979) (1982) (1982) (1982) (1985) (1987) (1988) (1989) r r r r r r r r r r r r r r

ko[OH-]°-36[03]2

-lO.Si 11.5

= k,[0H-][03] + k2[OH-]"-^[03]

= ko[OH-]°-5[03]

= ko[OH-]°-^[03]

= ko[OH-][03]

= kitOH-]°-28[03]1-5 + k2[OH-][03]

ko[OH-]°-12[03]

ko[OH-]°-«[03]2

4200 [OH'lCOj] in mol./min.

2880 [OH'lCOj] in mol./min.

kiCOH-jCOj] + k2[OH-][03]2

ki[03] + k2[OH-]1/2[03]3/2

3ki[OH-][03] + k2[OH-][03]2

3k,[OH-]C03] + k2[OH-][03]2

^ Extracted from Gurol & Singer (1982)

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-t—^W,-<5rrli»"

aqueous ozone reacts with dissolved substances via two pathways: (a) an indirect (radical) reaction pathway and (b) a direct reac¬

tion pathway.

The indirect reaction pathway involves the reaction of dissolved substances (solutes) with radicals (especially OH radi¬ cals) formed during ozone decomposition. Hoigne and co-workers in

various studies (1976,1982,1983) have reported that the indirect reaction pathway is initiated at a rate proportional to the

hydroxide (OH") ion concentration, generating two types of radi¬

cals: superoxide radical (O^) and hydroperoxy radical (HOg) , the protonated form of superoxide radical (pK^ =4.8). It is worth¬ while noting that one molecule of O^' radical and one molecule of

HO2 radical are formed per molecule of ozone consumed in the

initiation reaction. Each superoxide radical transfers an electron to an ozone molecule in a highly-selective reaction to form an

ozonide anion. The ozonide anion immediately decomposes through a series of reactions (see Figure 2.1 and Table 2.3) to a highly reactive (non-selective) hydroxyl radical. In pure water, the hydroxyl radical reacts with another molecule of ozone to regener¬

ate the superoxide radical which propagates the above-described

sequence of chain reactions (see reactions (2) through (5) in Table 2.3 and Figure 2.1). However, the sequence of chain reactions could be modified depending on the relative amounts of promoters, compounds that convert the hydroxyl radicals to ozone-selective

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03

L«-)o*>r*r

OH"

}*ͣ

(D

^0;

OJ^SHHO,* 1^---

ͨ

Oj

3*

•o-HO3—-^—1 OH I M-^i-< ^

0 {°i?"}^VTrt-Hoi

03 Oj+Ol+H^O

02

Figure 2.1 Reactions of dissolved ozone in pure water

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Table 2.3 Fundamental reactions of ozone decomposition (after Staehelin and Hoigne (1982,1985))

Initiation Rate constant

O3 + OH' ---ͨ HOj + 0{ ^^1 = 70 M'^s''' (1)

HO2 trz} O2" + H* K = 10-4.8 (!•)

Propagation

O2' + O3 ---ͨ O3" + O2 ^ = 1.6 X 10* M'''s''' (2) 0{ + H* ---ͨ HO3 ^5 = 5.2 X 10^° K'^s'^ (3) HO3---ͨ OH + Oj ^4 = 1.1 X 10^ s'^ (4)

Scavenging/teriuination

OH + O3---ͨ HO4---ͨ HO2 + O2 ^ = 3 X 10' M'^s''' (5)

OH + CO32---

ͨ

OH' + CO3'

^.6 = 4.2 X 10^ M'^'s"'' (6)

OH + HCO3'---ͨ OH' + HCO3 ^^.7 = 1.5 X 10^ M'''s''' (7)

OH + HPO^^'---

ͨ

OH' + HPO^'

^.8 = 5 X 10* M'''s*'' (8)

OH + H2P0^'---ͨ OH' + HgPO^ ^.9 = 2.2 X 10* M'^S"^ (9)

OH + PO^^'--->

ͣ

OH" + P0^2'

^,10 = 0.9 X 10* M'''s''' (10)

HO^ + HO^---ͨ HjOg + 2O3 (11)

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"iBTTi^H^sr'

radical scavengers (e.g. carbonate, bicarbonate, etc.). The

radical scavengers terminate the chain reactions by reacting with

non-selective radicals, for example the hydroxyl radical, to form

either stable compounds or radicals that do not promote the chain

reactions (see reactions (6) through (12) in Table 2.3 and Figure

2.1).

The direct reaction pathway involves reaction of ozone

molecules with dissolved substances. Hoigne and co-workers

(1982,1983) reported that this pathway is slow and selective. They

also reported that this pathway is favored at low pH and/or in the

presence of radical(OH) scavengers like carbonate and bicarbonate

ions. The relative importance of these reaction pathways during

ozone decay would not only depend on pH but also on solution

conditions.

Unfortunately, many of the earlier studies were carried

out under different solution conditions (i.e. different ionic

strengths, in the presence or absence of buffers, different types

and grades of chemicals, different initial ozone concentrations,

presence of scavengers and promoters, etc.) resulting in

system-specific rate constants and rate expressions. This could be one of

the reasons for the substantial differences among the findings of

various researchers. It is worth mentioning that in this study

different ozone decomposition rates were observed in experiments

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One other reason for the discrepancies could be due to

the use of different analytical techniques to measure the residual

dissolved ozone concentration. It is known (Gordon and Grumwell,

1983) that the iodometric method is subject to interference by decomposition products which may also be oxidants. Hence, an overprediction of the residual dissolved ozone concentration results in apparent slower decomposition kinetics. On the other hand, UV measurements of ozone are less sensitive to low ozone concentrations and there is considerable uncertainty (Hart et al., 1983) with regard to the wavelength and molar absorptivity values that should be used for measuring the dissolved ozone concen¬

tration. Inaccurate measurements can lead to erroneous kinetic

models for ozone decomposition.

The experimental conditions listed in Table 2.1 show that different studies have been conducted at different pH values and for different temperature ranges. It should be noted that activation energy values ranging from 40 kJ/mole to 115 kJ/mole have been reported in the literature (Sotelo et al., 1987) for ozone decomposition. Thus, it is clear from the above discussion that different experimental conditions can lead to kinetic models which are valid only under those conditions.

(ii) Effect of humic material

Staehelin and Hoigne (1985) proposed that various

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OxODBtBd Humic SubstflDce

Osxke Caisumed by Radical Process

OH* O3 "o:

Humic Substaoce

Ozooe CoQsuined

O3 by Direct Process

(HS) = JlctiyesiteM^lnmiciiiBteiiBlthatactsasa

^ ' diredrBactar, mmaor, promoter, 01or scavenger

Figure 2.2 Reactions of dissolved ozone with humic material

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radicals (direct reactors), some reacting directly producing

hydroxyl (0H-) radicals (initiators), and some promoting the

radical chain through reactions that produce species such as the

superoxide radical (Og"). Figure 2.2 summarizes these pathways.

Reckhow et al., (1986) and Legube and co-workers (1989,1991) have shown that humic material acts as a promoter and

possibly as an initiator of radical reactions. However, there is general disagreement among researchers as to whether or not humic

material acts as a radical scavenger. Staehelin and Hoigne (1985) proposed that humic material acts as a radical scavenger, but Feng and Legube (1991) recently reported the opposite. These con¬ flicting results could be due to the use of humic material from different sources. Reckhow et al. (1986) and Legvibe and co-workers (1989,1991) conducted their investigations using humic material extracted from natural waters (different water sources) by XAD hydrophobic resins. It should be noted that the characteris¬ tics of natural humic material depend mainly on the age and the

source. Staehelin and Hoigne (1985) and Hasten (1990) used

commercially-available humic material (Fluka humic acid) in their studies. It has been demonstrated by Malcolm and McCarthy (1986)

that commercial humic material consists of different functional

groups and elemental composition than that of aquatic natural humic material. Thus commercially-available humic material may not ade¬ quately represent the behavior of aquatic natural humic material.

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Hoigne 1985, Hasten 1990) has been used to compare results of various investigations.

Though there is disagreement among researchers on the

role of humic material in the radical reactions of ozone decompo¬

sition, many investigations have demonstrated that ozone decompo¬

sition follows first-order kinetics in both natural water and

synthetic waters containing humic material (Staehelin and Hoigne,

1985, Yurteri and Gurol, 1988, Legube et al., 1989). These investigators also report that the ozone depletion rate increases with increasing concentration of humic material. A few studies

have reported that the apparent first-order reaction kinetics is

preceded by an initial ozone demand due to rapid utilization of ozone during the first minute ( Reckhow et al., 1986 and Legxibe et al., 1989). The effect of initial ozone dose on the rate of depletion of dissolved ozone has been studied by a few investi¬ gators. Anderson et al. (1986) reported that in solutions

containing similar molar ratios of initial ozone dose and dissolved

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B. Kinetic models for aqueous ozone decomposition

(i) Ozone decomposition in pure water

Hoigne and colleagues (1982,1985) proposed that aqueous

ozone decomposition in the presence of dissolved substances could

be represented by the fundamental reactions listed in Table 2.3 (also see Figure 2.1). In this scheme, the initiation step is

characterized by an oxygen radical transfer from ozone to hydroxide

ion.

It is clear from the values of the rate constants that

the free-radical initiating step constitutes the rate-determining

step in the scheme. It is important to note that the generation of

one mole of superoxide radical ion O^ or its protonated form HOg

from the hydroxyl radical OH (reactions (!•) through (5) listed in

Table 2.3) implies that 1 mole of ozone is consumed. Hence all

species consuming hydroxyl radicals without regenerating superoxide radical ion (reactions (6) through (10) listed in Table 2.3) will

stabilize ozone in water.

Yurteri and Gurol (1988) considered reactions (1) through

(10) (see Table 2.3) and derived a mixed-order model for ozone

decomposition (see Appendix B for a complete derivation of the mixed-order model for pure water systems). The steady-state equations for the concentration of hydroxyl radical and superoxide radical are given as

2ki[0H"] [0,1

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where

Ekj ,[S,] = ks^JC032-] + k3 7[HC03-] + kggEHPO.^"] +

3--kj^^CH^PO,-] + ks^iotPO,-]

[O,"] = 2ki[0H-] 1 + ksCOj]

(2.2)

(2.3)

Using Equations (2.1) through (2.3), the mixed-order model can be

written as

^ =k^[03] ^y^'^io,f

(2.4)

where

-U-i

k,' = 3k,[OH'] with k, = 70 K'^s

.^4k,k5[OH-3 ^.^^ k. = 3 X 10' M-^s-l

Ikg.[Sj] = contribution from scavengers (defined in

Equation (2.2))

When the magnitude of the contribution from scavengers is large (for example in carbonate-buffered systems such as high alkalinity

waters), kg' becomes small. In such situations, the mixed-order

model (2.1) reduces to a simple first-order model given by

(2.5)

Staehelin and Hoigne (1982) arrived at the first-order expression based on theoretical and experimental observations.

(30)

where

-_^=k^[OH-][03]+k'2[OH-][03]2 (2.6)

at

k', = k^H = 111 M-1S-1

2koHjC8 k 3 X 10' M-1s-1

2 ks[S] 8

with kgCS] = k^CCOj^-]

and proposed the reaction mechanism listed in Table 2.4.

This mechanism involves a two-electron transfer process

or an oxygen atom transfer from O3 to OH ion. Here again, they

showed that in carbonate-buffered systems, the rate eguation can be represented by a simple first-order model. It is worth noting that

Tomiyasu et al. (1985) obtained k, = 40 M'^s"' value by conducting

experiments at pH 12.

(ii) Aqueous Ozone decomposition in the presence of humic

material

The decomposition mechanism of ozone is significantly modified by the presence of dissolved humic material. Staehelin and Hoigne (1985) proposed the following additional reactions to go

with the earlier mechanism listed in Table 2.3, when the humic

material acts as direct reactors, initiators, propagators and

scavengers:

O3 + D---ͨ products (direct reaction k^) (13)

(31)

Table 2.4 Fundamental reactions of ozone decomposition (after Tomiyasu, Fukutomi and Gordon (1985))

Initiation Rate constant

O3 + OH" ---ͨ H02' + O2

k, = 40 M"''s"'

(1)

Propagation/termination

O3 + H02' ---ͨ HO2 + Oj-

k2 = 2.2 X 10* M'^s"^

(2) HO2 t:=i O2' + H*

k^ = 10"*-®

(3) 0{ + O3 ---. O3- + O2

k^ = 1.6 X 10^ M"^s'

(4)

O3" + H2O ---ͨ OH + O2 + OH"

k5 = 20 - 30 s"''

(5)

03" + OH ---ͨ HO2 + 0{

k^ = 6 X 10' M"''s"''

(6) O3" + OH ---ͨ O3 + OH"

ky = 2.5 X 10' M"''s"''

(7) OH + O3 ——ͨ HO^--->ͣ HO2 + Og

kg = 3 X 10' M'^ls"'

(8)

OH + COj^'

---ͨ OH" + 003"

kg, = 4.2 X 10^ M"^S"'

(9)

003" + O3 ---ͨ Products (CO2 4 • O2 + 02") (10)

(32)

Yurteri and Gurol (1988) used reactions (11) through (16) in

addition to reactions (1) through (10) of the Staehelin and Hoigne (1985) mechanism (see Table 2.3) to derive an expression for ozone

decomposition in the presence of humic material and other solutes (see Appendix C for a complete derivation of the model equation).

The model gives

--^^^ = ko[D] + k,[I] + 3k,[0H-] +

(2ki[0H'] + ki[I]) , ,

^' ----^^—^ (kp[P] + 2k5[03]) (2.7)

It should be noted that Equation (2.7), describing ozone depletion in water containing dissolved humic material, is mixed-order (both

first and second-order) with respect to ozone. This generalized mixed-order rate equation can be simplified by assuming kp[P] »

SkgCOj] and fairly large scavenger concentrations and rewritten as

d[03]

--^ = w[03] (2.8)

where the specific utilization rate, w [time"''], is defined as

w = w^ + w„ + w*0 0

W^ = k„[D] + k,[I] W„ = 3k,[0H-]

W* = (2k,[0H-] + k,[I])kp[P]/Iks,[S,]

(33)

correlation for w based on these samples (R^ = 0.83) which is

given by

log(w) = -3.98 + 0.66pH + 0.611og(T0C) - 0.421og(alk./10) (2.9)

This expression successfully predicted ozone decomposition rates for 11 other natural water samples. The predicted values for w,

when substituted into Equation (2.9), corresponded well with the apparent first-order rate constants measured for these water samples. However, it should be noted that this expression does not include the effect of initial ozone dose, and does not account for

(34)

3. EXPERIMENTAL METHODS

A. INTRODUCTION

The experimental conditions were chosen to simulate the natural water characteristics of pH, alkalinity and total organic carbon content. In order to minimize the number of experiments required to evaluate the effect of various experimental conditions

on ozone decomposition, a factorial design approach was considered

for the experiments. Because it was difficult to elucidate the effects of individual parameters in a single experimental set in the factorial approach, it was decided to conduct the experiments in batches. Each batch consisted of a series of experiments, for example varying the pH of the water and studying the effect of pH

on decomposition of ozone while holding all other constituents

constant. This approach, though not optimal from a statistical

standpoint, provided an opportunity for studying the individual

effects of each variable, so that the experimental results could be compared to earlier investigations and used to examine various

mechanisms proposed for ozone decomposition. All the experiments

conducted during the study are listed in Table 3.1. B. EXPERIMENTAL SET-UP

(35)

Table 3.1 Experimental Matrix

Temperature = 22 "C and Total phosphate = 1 x 10'^ M

pH = 5.5

Total Carbonate, 10'^ Humic Acid, mg/L 0.0 0.0

2.0 0.0 - 4.0 pH = 6.5

Total Carbonate, 10'^ Humic Acid, mg/L

0.0 0.0

2.0, 8.0 0.0 pH = 7.0

Total Carbonate, lO'^M Humic Acid, mg/L

0.0, 1.0 0.0

2.0 0.0 - 4.0

8.0, 10.0 0.0 pH = 8.0

Total Carbonate, 10"^ Humic Acid, mg/L

0.0 0.0

2.0, 8.0 0.0 pH = 8.5

Total Carbonate, 10"^ Humic Acid, mg/L

0.0 - 1.0 0.0

2.0 0.0

(36)

ͣ

'^^^ESr^rr^f^^r^

generator (W.R. Grace & Co., Model LG-2-L1) from a standard gas

cylinder through 1/4" Tygon plastic tubing. All tubing from the

ozone generator to the gas traps and exhaust were 1/8" Teflon

tubing. To avoid corrosion, all fittings and valves in this set-up

were either stainless steel or Teflon.

The ozone-oxygen mixture from the ozone generator was

bubbled into deionized distilled water through a sintered-glass

dispersion tube at the bottom of a glass carboy. The 5-gallon

glass carboy was covered at the top with a specially-designed glass

cap, gasketted with a Teflon ring and held in place by a horseshoe-shaped spring clamp to ensure a tight seal. The glass cap had four ports: for ozone-oxygen mixture inlet, ozone-oxygen mixture outlet, distilled water inlet and ozonated water outlet. The water in the glass carboy was stirred by a Teflon-coated magnetic stirring bar. The off-gas from the carboy was bubbled through a series of traps containing saturated sodium hydroxide solution to strip the maximum amount of ozone from the gas stream before venting it to the atmosphere through a negative-draft exhaust system. Figure 3.1 depicts the above assembly.

The experimental batch reactor consisted of a 200-mL

graduated glass cylinder with a sampling port at the bottom from

which samples could be withdrawn for the measurement of residual

dissolved ozone. The reactor was provided with a floating Teflon cover in order to maintain the test solution headspace-free, regardless of the volume of sample withdrawn for ozone analysis.

(37)

Ozonek

To Vent T

1

_II—

Generator 1

Dessicant

1 r

1,

A

e

X

O _Dei

---1^^

onized

^

/^

J

1^

Glass

Carboy

S

^

IT-o

3

O

0 o ooo

J

Distilled

Water Inlet

Ozonated

Water Ozone

Traps

(38)

solution due to volatilization. The contents of the reactor were well-mixed by a Teflon-coated magnetic stirring bar. To keep the temperature of the contents of the reactor constant, a water jacket was provided around the reactor. All the experiments were conducted at one temperature 22 ± I'C. Because there was almost no fluctuation in this temperature, the water jacket was not required. Figure 3.2 shows the batch reactor configuration.

C. EXPERIMENTAL PROCEDURE

The experimental investigation was conducted in three phases. In the first phase, a study of the effect of pH was con¬ ducted, followed by modelling of the ozone decomposition rate and determination of the characteristic kinetic parameters. In the

second phase, the effect of total inorganic carbon (C^) on ozone

decomposition was studied at various pH values. In the third phase, the effect of organic carbon (Aldrich Humic Acid) on the ozone decomposition rate was studied at various pH values and inorganic carbon content. Finally, ozone depletion in natural waters was tested by conducting batch kinetic experiments on East Bay Municipal Utilities District, CA (Sobrante Water Treatment

Plant) settled water and Hackensack, NJ raw water.

(i) INITIAL PREPARATION

(39)

FtomCoMtant

TBtspontm Wata*

B^

^---Sample Port

llwnnooifBtBr

'TefkaCap

Reactor

StirBi

<z

ToWeta-Beth

(40)

Particulate Air (HEPA) filter) through the solution for about 4 to

6 hours. The air outlet was connected to a sodium hydroxide ozone trap. The ozone-demand-free water thus prepared was used to make all reagents necessary for the experiments. In order to remove any

ozone-demanding impurities from glassware, all the glassware used

in the experiments was rinsed first with distilled water and then

with ozone-demand-free water, and dried before use.

(ii) Test Solution Preparation and Seuapling Methods

The test solutions were prepared using ozone-demand-free

water. The pH and buffer concentration of the solutions were

adjusted using pre-determined amounts of phosphate buffer (KHgPO^

/ KjHPO^ (analytical reagent grade, British Drug House (BDH) Chemi¬

cals, Poole, England). The desired pH was achieved by adding small volumes of dilute HCl/NaOH (ACS-certified Fisher Scientific, Fair Lawn, NJ), while the contents of the reactor were mixed with a magnetic stirring bar. The pH was measured using an Accumet 915 pH meter (Fisher Scientific, Pittsburgh, Pa) which had a temperature

probe to compensate for temperature changes. The pH meter was

calibrated daily with standard buffers ('Baker analyzed', Fisher

Scientific Chemical Co., Fair Lawn, NJ) . All experiments were studied in solutions of 0.01 M total phosphate species concentra¬ tion (Py). Most of the experiments were conducted at pH 5.5, 6.5, 7.0, 8.0 and 8.5. However, at the start of the study, experiments at pH 2.0 were conducted to make sure that there was no loss of

(41)

beginning and end of each experiment, pH and temperature were

measured.

In order to study the effect of inorganic carbon (HCO^'and

COj^") on aqueous ozone decomposition, the total carbonate ion

content of the test solution was varied from 0 to 10 x 10"^ M at the

various pH values mentioned above. A solution of 0.1 M sodium carbonate (analytical reagent grade, ACS-certified, Mallinckrodt Chemical Co., Paris, KY) was used to adjust the total inorganic carbon content. It should be noted that the stirring speed was lowered after the addition of sodium carbonate to avoid degassing

of carbon dioxide.

The effect of natural organic material on aqueous ozone

decomposition was studied by adding pre-determined amounts of

commercial humic acid (Aldrich Chemical Co., Milwaukee, WI). Humic acid solution was prepared by dissolving a known amount of hvimic material in 0.1 N sodium hydroxide (ACS-certified, Fisher Scientif¬ ic Chemical Co., Fair Lawn, NJ) . The humic acid (HA) concentration was varied from 0 to 4 mg/L. The total organic carbon (TOC) content of the Aldrich humic acid was 31% by weight. This set of

experiments was conducted at pH 5.5 and 7.0, and the C^ content

was varied from 0 to 2 x 10'^ M.

Once the contents of the reactor were adjusted to the

desired pH, C^ and TOC levels, pre-calculated amounts of dissolved

ozone from stock solutions containing up to 30 mg/L dissolved ozone

(42)

b!!ll.;«K*f,r. •

initial ozone concentration of 3.0 mg/L was used. Some experiments were also conducted with 1.0 and 0.3 mg/L as initial ozone concentrations.

Immediately after the ozone was added to the test solution, the contents of the reactor were covered with a Teflon float and stirred with the magnetic stirrer. The first sample was withdrawn from the reactor within one minute of ozone addition.

This was done to assure complete mixing of the reactor contents before sampling. The first few milliliters from the sampling port

were discarded and then 20 mL samples were quickly withdrawn into a clean graduate cylinder at desired sampling time intervals. The 20 mL samples were immediately added to 20 mL of indigo solution (see below) to measure the residual dissolved ozone concentration.

The same experimental procedure described above was used for

natural water samples. As mentioned earlier natural water samples

were obtained from East Bay Municipal Utilities District, CA

(Sobrante Water Treatment Plant, settled water) and Hackensack, NJ (raw water). These water samples were kept refrigerated (10 "C cold rooms) and within two days of receipt of the samples batch kinetic experiments were conducted on them. The experiments were

conducted without any dilution of the sample. Only small quanti¬

(43)

D. ANALYTICAL PROCEDURES FOR DISSOLVED OZONE MEASUREMENT

(i) lodometric Method: The aqueous ozone concentration in the stock solution was measured by this method (Gurol, 1980). Fifty mL

of ozonated water was withdrawn in a volumetric pipette, washed

previously with ozonated water, after discarding the first few

milliliters of solution from the outlet of the glass carboy. The

sample was immediately transferred to a beaker containing 40 ml of 2% (w/w) potassium iodide (analytical reagent grade, ACS-certified,

EM Science, Gibbstown, NJ) . The pH of the resultant mixture was

reduced to a value below 2 by adding a few milliliters of concen¬

trated sulfuric acid (ACS-certified, Fisher Scientific, Fair Lawn, NJ) to transform the iodate formed to iodine. The solution was titrated with 0.005 N sodium thiosulfate (analytical reagent grade,

ACS-certified, EM Science, Gibbstown, NJ) using starch as an

indicator. The endpoint of the titration was the disappearance of the dark blue color. The sodium thiosulfate was standardized with potassium dichromate (ACS-certified, Fisher Scientific, Fair Lawn, NJ) (Standard Methods, 1989). The relevant reactions are

Absorption:

O3 + 2H* + 21" ---> I2 + HgO + Oj (3.1) Titration:

2S2O32- + I2 ---> 21' + S^0^2- (3.2)

Thus, two moles of sodium thiosulfate are equivalent to one mole of

(44)

relationship:

mg of ozone ^ 24NV, ^

L of solution V2

where N = Normality of the titrant (sodium thiosulfate) V, = mL of titrant added to reach endpoint

Vg = sample voliime in liters

The titration was repeated at least twice before the start of each experiment and the average titrant volvime was used to calculate the

stock ozone concentration. It is worthwhile to note that the

dissolved ozone concentration measured by the iodometric method may be inaccurate due to the potential presence of other oxidant species (e.g. ozone decomposition by-products) in the stock solution. Hence, the stock solution concentration was also

measured by the UV method.

(ii) UV Method: A quartz cuvette (1 cm) was cleaned and filled

with ozone-demand-free water. Its absorbance was measured and set

to zero at a wavelength of 258 nm in a Spectronic 1201 spectropho¬ tometer (Milton Roy Co., Rochester, NY) . The cuvette was then rinsed with the stock ozone solution and filled completely. The cuvette was immediately covered with a Teflon cap allowing no head-space for volatilization. The UV absorbance of the sample was measured at 258 nm. This procedure was repeated a number of times

(45)

following formula was used to calculate the ozone concentration:

mg of ozone ^ a* (change in absorbance) (3.4)

L of solution

where

a = molar absorptivity of ozone

a = 2950 L/(mole cm) (Standard Methods, 1989)

It is worthwhile noting that there is some uncertainty in the

literature (Hart et al., 1983; Gordon and Grumwell, 1983) with regard to the molar absorptivity of ozone and the wavelength at

which the absorbance should be measured. However, Standard Methods

(1989) suggests the use of 2950 L/(mole cm) as the molar absorpti¬ vity at 258 nm wavelength for the UV method (these values were used for measurements and calculations in this study). The ozone concentration calculated using the UV method was usually found to be slightly lower (1%) than the value calculated using the

iodomet-ric method. This discrepancy suggests that the iodometiodomet-ric method

might be susceptible to interference from other oxidants formed during ozone decomposition or that the value of the UV molar absorptivity might be lower than that used in Equation (3.4).

Because there are no other reliable standard methods available to

measure dissolved ozone concentration, the average of the dissolved ozone concentrations measured by the iodometric was used in subse¬

quent calculations (e.g., as a standard for the indigo method).

(46)

the measurements were repeated with freshly-prepared reagents and

cleaned glassware.

(iii) Indigo Method: The residual dissolved ozone concentration in the kinetic experiments was followed by the indigo method which is based on the property of the ozone molecule to oxidize the blue-colored indigo dye (Bader and Hoigne, 1981). This method is simple, quick and sensitive to low concentrations of ozone. It is

found to be insensitive to interference from other oxidants that

may be products of ozone decomposition.

A solution of the trisulfonated potassium salt of indigo

(C^^H^KjNgO^^Sj, Molecular weight = 616.7 (Aldrich Chemical Co.,

Milwaukee, WI)) was prepared in ozone-demand-free water. A mixture

of HjPO^ (analytical reagent grade, ACS-certified, EM Science,

Gibbstown, NJ) and KjHjPO^ (analytical reagent grade, British Drug

House, Poole, England) was added so that the pH of the resultant mixture would be in the favorable region of 2 to 4 (Standard

Methods, 1989). In this pH range, self-decomposition of ozone is minimal. The indigo stock solution thus prepared was diluted 100

times and its absorbance was measured in a Spectronic 1201

spectrophotometer (Milton Roy Co., Rochester, NY) at a wavelength of 600 nm in a 1-cm glass cuvette. Indigo solution has a maximum absorbance at this wavelength, with a molar absorptivity of 1.44 x

10* L/(mole-cm) . Before each use, the absorbance of the indigo was checked and if it fell below 95% of the initial value the stock

(47)

Ozone reacts with the single C=C double bond in the indigo molecule, decreasing the absorbance of the indigo solution. In the pH range 2 to 4, the decrease in absorbance (AA) is proportional to the concentration of ozone. For purposes of calibration, stock solutions containing known concentrations of ozone (mean value measured using the iodometric method) were added to indigo solution and the change in absorbance was measured. The calibration curve thus obtained was used to measure the residual dissolved ozone concentrations in the unknown sample generated during the kinetic experiments. Figure 3.3 is an example calibra¬

tion curve.

The slope of the straight line correlating ozone concentration to AA is a function of the initial concentration of the indigo solution. Thus, the sensitivity of the indigo method can be adjusted by changing the initial concentration of the indigo

solution. In this study, three different indigo solutions were used to measure dissolved ozone concentration (for ranges 0 to 6,

0 to 2 and 0 to 1 mg of dissolved ozone/L). These solutions were

prepared by diluting the stock solution of indigo in various ratios such that the final absorbance of the indigo solutions, hence AA, could be easily measured by the spectrophotometer. The corre¬

sponding calibration curves are shown in Figures 3.3 to 3.5. The

differences in the slopes of these calibration curves should be noted.

(48)

the iodometric method. As noted earlier, the iodometric method is subject to interference due to ozone decomposition products. Such interferences were minimized by conducting the iodometric titrati¬ ons at a pH value below 2. The UV method was used to check the measurements made by the iodometric procedure. If the average dis¬ solved ozone concentration measured by the iodometric method was not within 1% of the value measured by the UV method, the measure¬ ments were repeated with freshly-prepared reagents and cleaned glassware. Whenever new indigo solutions were prepared, they were calibrated before use and their corresponding slopes were used to

(49)

0.500

0.400 —

< <1

0)

o

o 0.300

JQ O

w

D

.E 0.200 —

O)

c o

u

0.100

0.000

Conditions X = 600 nm

b (cell constant) = 1 cm [Indigo] = 125 yu,M

0 2 4 6 8

Concentration of ozone, C mg/L

Figure 3.3 Calibration Curve I of indigo method for 0-6 mg/L

(50)

0.200

< <d

0)

o

c D XI O

V)

D

C

D

x:

o

0.150 —

0.100

0.050

0.000

Conditions

A = 600 nm

b (cell constant) = 1 cm

[Indigo] = 50 ^lU

Concentration of ozone, C mg/L

Figure 3.4 Calibration Curve II of indigo method for 0 - 2 mg/L

dissolved ozone.

(51)

0.050

0.040

< <]

0) o

o 0.030

-Q

V-O

(/)

-Q D

.E 0.020

Q) c

O

U

0.010

0.000

Conditions X = 600 nm

b (cell constant) = 1 cm [Indigo] = 25 fiM

0.000 0.200 0.400 0.600 0.800 1.000

Concentration of ozone, C mg/L

Figure 3.5 Calibration Curve III of indigo method for 0-1 mg/L

(52)

4. RESULTS AND DISCUSSION

A. INTRODUCTION

Before studying the effect of pH on aqueous ozone

decomposition, experiments were conducted to verify whether there

was any ozone loss due to volatilization during transfer of ozone

stock solution and from the experimental batch reactor. This

verification was carried out by conducting experiments at pH 2.0.

At this pH, the decomposition rate of ozone is minimal because the

concentration of hydroxide ion (OH") which initiates ozone decomp¬

osition in water is very small. Thus, any decrease in aqueous

ozone concentration could be wholly attributed to ozone volatiliza¬

tion from the reactor. Figure 4.1 shows that there was no change

in ozone concentration in the reactor which proves that there was

no volatilization of ozone. This experiment also shows that there

was no loss of ozone during the transfer of ozone stock solution to

the batch reactor since the calculated initial ozone concentration

was detected in the solution.

B. EFFECT OF pH

The effect of pH on ozone decomposition was studied at pH

5.5, 6.5, 7.0, 8.0 and 8.5. These pH values were selected so as to

simulate the pH range of a variety of natural waters. Phosphate

buffers (KHjPO^ and KjHPO^) were used in all experiments

(53)

CD E

c o

N

O

c

o

c CD O c

o

O

1 —

0

Phosphate Buffer

Pj = 0.01 M

o

T = 22 C

C = 3 mg/L

0 10 20 30 40 50 60 70

Time (minutes)

80 90 100

Figure 4.1 Decomposition of ozone in water at pH 2.0,

(54)

(total phosphate species concentration) and control the pH. All

the experiments were carried out at 22 ± 1 °C. The experimental

procedure section in Chapter 3 gives a more detailed and complete

description of the experimental conditions and methods.

In order to analyze the experimental data, aqueous ozone

decomposition at various pH values was modeled by mixed-order (a

combination of first and second-order rate terms), first-order and

second-order rate equations. The mixed-order model reported by

Tomiyasu et al. (1985), Yurteri and Gurol (1988) and Grasso and

Weber (1989) has the form

(4.1)

where,

with

and

k,' = 3k, [OH"]

k, = 70 M'^s"''

(4.2)

based on Staehelin and Hoigne (1982)

k',=_4k,k5[OH"] 1^-1

with kg = 3 X 10'' M'^s (4.3)

Ikg ,-[Sj] = contribution from scavengers.

The mixed-order model was developed based essentially on the

mechanism proposed by Staehelin and Hoigne (1982). A detailed

discussion on the mechanism and the fundamental rate equations

(55)

It should be noted that if the kj' values (in Equation 4.1) are

very small, then the mixed-order model reduces to a simple

first-order model.

In order to evaluate the applicability of the mixed-order

model, the k', and k'g values were calculated using the experimental

data. In order to solve the differential equation (Eqution 4.1) a

FORTRAN program was written using the modified Newton's algorithm

(Vandergraft, 1986). But the roots of the differential equation

suggested that there were no unique k', and k'g values i.e. the root

domain contained multiple solutions. Reduction in increment size in

the algorithm lead to divergence. Therfore, a quasi-theoretical

approach was used to evaluate the kg' values. Equation (4.1) was

integrated by the method of partial fractions with the initial

condition

[O3] = [03](, at t = 0 (4.4)

to give

[O3] = ---±--- (4.5)

1 . ^'z ,, .H^'it,

[Oslo k^

d-e" '')

Equation (4.5) can be rearranged as

k/2 = _Z^_ [^^-JLJL] (4.6)

' , icV ,, [O3] [03]o^ ^ ^

(e 1 - 1)

(56)

concentration in these experiments. The pH maintained in these

experiments and the fundamental rate constant (k,) listed along

with Equation (4.2) were also used in the calculations (the k,

value based on Staehelin and Hoigne (1982) was used. The k, value

based on Tomiyasu et al. (1985) (see Chapter 2) was not used in the

calculation since that value was obtained from experiments

conducted at pH 12 which is significantly different from the

present experimental conditions. (Example calculations are

illustrated in Appendix D.) The arithmetic mean of these calculat¬

ed k'j values were used to compare with the corresponding theoret¬

ical values. It is important to note that the k'g values varied

with time randomly (see Appendix D for further details). This

might be due to lopsided distortion of the slight error in low

residual ozone concentration measurements by the non-linear

functions in Equation 4.6.

The theoretical k'g values were calculated using Equation

(4.3). The total scavenger contribution (Sk^ ,8,.) was evaluated

using the phosphate species concentrations along with the appropri¬

ate fundamental rate constants (listed in Table 2.3 in Chapter 2).

Phosphate ion (HjPO^", HPO^^', and PO^^') concentrations were calculat¬

ed using the corresponding dissociation fractions (which are func¬

tions of pH and the acidity constants for phosphoric acid) and the

total phosphate species concentration (see example calculations in

(57)

The kg' values computed by the above methods are listed

in Table 4.1. Table 4.1 also lists the range over which the

experimental kg' values varied with time. As mentioned earlier,

the variation with time could be due to magnification of slight

error in low residual ozone concentration measurements by the

exponential functions in Equation 4.6. Comparing the experimental

and theoretical kg' values, it is clear that there is little agree¬

ment between these values. The experimental values (including the

range over which these values varied) are 5 to 500 times the

theoretical values. Staehelin and Hoigne (1985) observed similar

discrepancies, i.e. their experimental kg' values were much greater

than the theoretical kg' values. They reported that this discrep¬

ancy may be due to the presence of impurities. The impurities can

lead to initiation reactions that were unaccounted for.

Using the generalized mixed-order model, it can be shown

that the lack of agreement between the theoretical and experimental

kg' values reported here may be due to the presence of initiation

reactions that are unaccounted for. The generalized mixed-order

model (see Chapter 2 and Appendix C for a detailed derivation) can be written as

1 d[033

- ^

_{[03] dt}

ͣͣ

^' = ko[D] +kj[l] +3ki[0H-] +

2ki[0H'] + ki[I]

(58)

Table 4.1 Comparison of theoretical and experimental kg' in

mixed-order model for experiments^ at 22 "C with no inorganic &

organic carbon

PH

_{Theoretical (kg' xlO^)}

(L/(mg(min) (Hoigne^)

Experimental (kj'xlO^)

(L/(mg(min)) (Hoigne^) 5.5 6.5 7.0 8.0 8.5 0.015 0.125 0.320 2.224 6.790

3.320 ± 0.90

4.870 ± 1.10

6.686 ± 1.20 13.884 ± 2.40 29.915 ± 4.10

Phosphate-buffered with P^ = 0.01 M.

k, = 70 M"' s'^ (after Hoigne et al. (1982)).

Table 4.2 Summary of rate constants for first and second-order

models for experiments^ at 22'C with no inorganic and organic

_carbon

pH

First-order

K', 1/min

Second-order

K'2 L/(mg(min))

3k,[OH*] xlO^

1/min

(Hoigne^)

5.5 0.029

(r^ = 0.93)

6.5 0.051

(r^ = 0.99)

7.0 0.096

(r^ = 0.99)

8.0 0.153

(r^ = 0.99)

8.5 0.310

(r^ = 0.98)

0.030 (r^ = 0.97)

0.125 (r^ = 0.81)

0.371 (r^ = 0.85)

0.356 (r^ = 0.85)

2.908 (r^ = 0.89)

0.031

0.314

0.993

9.930

31.400

' Phosphate-buffered with P^ = 0.01 M.

(59)

For clean systems containing no direct reactors, such as the

aforementioned experimental test solution containing only phosphate

ions, Equation (4.7) can be simplified by assuming that 2k5[03] »

kp[P] and rewritten as

1 dro,]

^ •• ^

ͣ

' = k, [I] + 3k, [OH'] +

[03] dt

2ki[0H"] + k,[I]

2k5[03] (4.8)

Equation (4.8) can be rewritten as

-£^ =k^[03] *k^[03]2

(4.9)

where

rff =

k"i = 3ki[0H-] + kj[I] (4.10)

and

k^ =

2ki[0H"] + k,[l]

2JCs.i[S,-] 2k5 (4.11)

Equations (4.2) and (4.3) could be substituted into Equations

(4.10) and (4.11) and rewritten as

r// = V/

(60)

K", =k'2 + ^^^'^^^ (4.13)

Equations (4.12) and (4.13) clearly show that the rate

constants (k"2 and k",) in Equation (4.9) will have higher values

than the rate constants (k'j and k'^) in Equation (4.1) if initi¬

ators are present. This could be one of the reasons for the lack

of agreement between the theoretical and experimental values of

k*2- It was beyond the scope of this study to separately evaluate

the contribution of initiators, other than OH" ions, to k," and kg".

These undefined initiation reactions can be avoided by using high

concentrations of a radical scavenger such as carbonate and bicar¬

bonate ions. This system will be discussed later in this Chapter.

Accordingly, the experimental data for various pH values

were modeled using empirical first-order and second-order rate

equations. Figure 4.2 depicts first-order ozone decomposition at

various pH values. Table 4.2 lists the corresponding first-order

rate constants and the theoretical 3k,[OH'] values for comparison

(based on Staehelin and Hoigne (1982)). Also listed are the rate

constants determined from a second-order fit of the data. It is

clear from the figure and regression coefficients (r^ values)

listed in the table that the first-order model fits the experi¬

mental data very well. Figure 4.2 shows that the rate of ozone

decomposition increased with increase in pH. This is due to the

increase in hydroxide ion concentration which initiates the ozone

decomposition cycle (see Figure 2.1 in Chapter 2). A number of

(61)

0

-2 —

o ^ -3

-4 —

-5

0 10 20 30 40

Time (minutes)

50 60

Figure 4.2 Effect of pH on ozone decomposition in water

(62)

reported the same observation. The extent of increase in the

first-order rate constant is proportional to the hydroxide ion

concentration ([OH']) in the test solutions. Figure 4.3 illus¬

trates this observation. Staehelin and Hoigne (1982) and other

researchers (see Table 2.2 in Chapter 2) have reported similar

results. It is important to note that the first-order rate

constant did not vary linearly with hydroxide ion concentration as

expected based on the k/ = 3k,[OH'] relationship. This behavior

could be due to other initiation reactions occuring due to the

presence of impurities in the test solution.

Figure 4.4 illustrates the second-order plots and Figure

4.5 depicts the predictions of the first-order and second-order

models. These figures clearly show the lack of fit of the second-order model.

In order to verify the first-order nature of aqueous ozone decomposition, the effect of initial ozone concentration was studied. Experiments were conducted at pH 7 with initial ozone

concentrations varying from 0.3 mg/L to 3 mg/L. The experimental

results illustrated in Figure 4.6 indicate that the first-order

ozone decomposition rate constant, represented by the slope of the line, is independent of initial ozone concentration. However, the slight deviation from first-order behavior at low residual ozone concentrations suggests a weak effect of the second-order rate con¬