EPA/600/3-85/031
April 1985
REACTIONS OF OZONE WITH ORGANICS IN AQUEOUS SOLUTIONS
by
C. H. Kuo
Mississippi State University
Mississippi State, Mississippi 39762
Cooperative Agreement CR808799
EPA Project Officer
H. M. Barnes
Atmospheric Chemistry and Physics Division
Atmospheric Sciences Research Laboratory
Research Triangle Park, North Carolina 27711
ATMOSPHERIC SCIENCES RESEARCH LABORATORY
OFFICE OF RESEARCH AND DEVELOPMENT
U. S. ENVIRONMENTAL PROTECTION AGENCY
RESEARCH TRIANGLE PARK, NORTH CAROLINA 27711
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TECHNICAL REPORT DATA
(Please read Inslrucno'U on the reverie icfore conpleting)
1. REPORT no
EPA/600/3-85/031
4. TITLE AND SUBTITLE
REACTIONS OF OZONE. WITH ORGANICS
IN AQUEOUS SOLUTIONS
3. REC
PBS5-191179
5. REPORT DATE
April 1985
6. PERFORMING ORGANIZATION CODE
7. AUTHOR(S)
C. H. Kuo
8. PERFORMING ORGANIZATION REPORT NO
9. PERFORMING ORGANIZATION NAME AND ADDRESS
Mississippi State University
Deparrtmer.t of Chemical Engineering
Mississippi State, Mississippi 39762
10. PROGRAM ELEMENT NO.
CDTA1D/02 Task 1585 FY-85
11. CONTRACT/GRANT NO.
CR 808799
12. SPONSORING AGENCY NAME AND ADDRESS
Atmospheric Sciences Research Laboratory - RTP, NC
Office of Research and Development
U.S. Environmental Protection Agency
Research Triangle Park, North Carolina 27711
13. TYPE OF RE PORT AND PERIOD COVE RED
Final
11/81-11/84
14. SPONSORING AGENCY CODE
EPA/600/09
15. SUPPLEMENTARY NOTES
16. ABSTRACT
Rates of ozonation of some organic pollutants in the aqueous phase were studiea.
Experiments were conducted in distilled water with pH varying between 2 and 7; the
solution temperature ranged from 5°C to 35°C.
Aromatic amines were very reactive with ozone and the reactions were second order.
The rate of the naphthylamine/ozone reaction increased with temperature. The
aniline/ozone reaction was faster, but was unaffected by temperature changes.
The toluene/ozone reaction exhibited a moderate rate. The order of ti.e reaction
changed with pH changes. PAHs including naphthalene, anthracene, and phenanthrene
reacted with ozone according to second order kinetics. The reaction between ozone
and anthracene was the fastest of all systems studied and the reaction rate was
nearly independent of acidity.
KEY WORDS AND DOCUMENT ANALYSIS
DESCRIPTORS
b.IDENTIFIERS/OPEN ENDED TERMS
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18. DISTRIBUTION STATEMENT
RELEASE TO PUBLIC
19. StCUR TY CLASS (Tins Kcporl)
UNCLASSIFIED
21. NO. OF PAGES
67
•JO. -ECUfilTY CLASS (Tins page I
UNCLASSIFIED
22. PRICE
EPA Form 2220-1 (R«». 4-77) PREVIOUS EDITION is OBSOLETE
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NOTICE
The information in this document has been funded wholly by the United
States Environmental Protection Agency under CR808799 to Mississippi State
University. It has been subject to the Agency's peer and administrative
review, and it has been approved for publication as an EPA document. Mention
of trade names or commercial products does not constitute endorsement or
recommendation for use.
ii
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ABSTRACT
Rates of ozonation of some aromatic pollutants in the aqueous phase were
studied by the stopped-flow technique. The kinetic experiments were conducted
in distilled water and in aqueous solutions of pH values ranging from 2 to 7
at 5 to 35°C,
Aromatic amines including aniline and a-naphthylamine were very reactive
with ozone and the ozonation reactions were second order. The rate of
a-naphthylamine-ozone reaction was enhanced by temperature, but the faster
reaction between aniline and ozone remained e.L a nearly constant rate for all
temperatures.
The order of reaction between toluene and ozone varied with acidity though
the rate of reaction was moderate. Polycyclic aromatic hydrocarbons including
naphthalene, phenanthrene and anthracene were reactive with ozone according to.
second order kinetics. Rates of the naphthalene-ozone and phenanthrene-ozone
reactions increased as the temperature increased and acidity of the solutions
decreased. The reaction between anthracene and ozone was the fastest among
the systems investigated, and the rate of reaction was nearly independent of
acidity.
ill
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CONTENTS
Abstract iii
Figures vi
Tables vii
Abbreviations and Symbols viii
Acknowledgement xi
1. Introduction 1
2. Conclusions 2
3. Recommendations . . 3
4. Materials and Equipment 4
Production of ozone gas 4
Preparation of buffer and aqueous solutions 5
Stopped-flow spectrophotoraeter system 6
Data acquisition system. , 9
5. Experimental Procedures and Analyses 11
Determinations of absorption coefficients 11
Selections of wavelengths 12
Selections of time constants 13
Kinetic experiments and analyses , 17
6. Results and Discussion 25
Ozonation of aromatic amines .....25
Ozonation of toluene 33
Ozonation of naphthalene. 41
Ozonation of phenanthrene .43
Ozonation of anthracene 47
References 52
Preceding page blank
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Figures
Number Page
1. Block diagram of stopped-flow spectrophotometer system 7
2. Absorbance of aniline in water at various wavelengths 14
3. Absorbance curves with different time constants... 16
4. Absorbance curves during early period of reaction 18
5. Absorbance changes during the aniline-ozone reaction 26
6. Determination of order with respect to ozone in the
aniline-ozone reaction, 27
7. Determination of order with respect to aniline in the
aniline-ozone reaction 29
8. Effect of temperature on the c'.-naphthylamine-ozone reac tion. .. .31
9. Absorbance changes during the toluene-ozone reaction 34
10. Determination of order with respect to toluene ...35
11. Effect of acidity on the toluene-ozone reaction ..37
12. Effect of temperature on the toluene-ozone reaction in
acidic solutions -9
13. Effect of temperature on the toluene-ozone reaction
in neutral solutions 40
x
14. Effect of acidity on the phenant'nrene-ozone reaction 45
15. Eff«:ct of temperature on the phenanthrene-ozone reaction 46
16. Effect of temperature on the anthracene ozone-reaction. 49
vi
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TABLES
Nurr'oer Page
1. Optimum wavelengths and time constants 19
2. Average rate constants for ozonation of aromatic
amines in water 30
3. Average rate constants for the toluene-ozone reaction 38
4. Average rate constants for the naphthalene-ozone reaction....A3
5. Average rate constants for the phenanthrene-ozone reaction...47
6. Average rate constants for the anthracene-ozone reaction 51
vii
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LIST OF ABBREVIATIONS AND SYMBOLS
ABBREVIATIONS
ac —alternating current
A/D —analog to digital
C •—degree centigrade
cm —centimeter
dc —direct current
kcal/g mole —kilocalorie per gram mole
1 —liter
M —molar or moles of solute per liter of solution
M cm --per molar per centimeter
l:~* s —per molar per second
nm —nanometer
pH ---logarithm of the reciprocal of hydrogen ion concentration
in gram atoms per liter
pKa —logarithm of the reciprocal of dissociation constant
psig —pound per square inch at gauge pressure
s —second
ms —millisecond
v —volt
vtii
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SYMBOLS
A(t) —absorbance as a function of time
b —length of light path
C^ —concentration of ozone
C^Q —initial concentration of ozone
£3 —concentration of organic reactant
CgQ —initial concentration of organic reactant
C^ —concentration of component i
HC1 —hydrogen chloride
H^O —water
HjPO^ —phosphoric acid
I —ionic strength
!„ —iodine
k. —rate constant for decomposition reaction
k~ —rate constant for ozonation reaction
k —reaction rate constant
k1 •—apparent rate constant
KI —potassium iodide
KOH •—potassium hydroxide
1 •—order of decomposition reaction of ozone
m —order with respect to ozone concentration in ozonation
reaction
n •—order with respect to concentration of organic reactant
in ozonation reaction
•—sodium dihydrogen phosphate
'—disodium hydrog-in phosphate
ix
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Nal — sodium iodide
NaOH — codium hydroxide
^2^203 — sodium thiosulfate
^^ — sodium tetrathionate
° — oxygen
— ozone
— reaction time
— voltage
— number of ion charges for component i
— absorption coefficient
— time constant
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ACKNOWLEDGEMENTS
The experimental work in this project was accomplished by Linda P.
Cornell, Steven C. Peng, Peter C. Wang, and Hanna M. Matta. The kinetic data
are documented in three M.S. Theses by Peng, Wang, and Matta, and in- a Ph. D.
Dissertation under preparation by Co-nell. Many valuable suggestions and
advises were offered by Dr. H. M. Barnes, the Project Officer from the U. S.
Environmental ptot-ec t_ ion Agency,
xi
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SECTION 1
INTRODUCTION
Aromatic compounds are among the major pollutants emitted into the
atmosphere from mobile and stationary sources such as automobiles, petroleum
refineries and chemical raanufacturing and fuel combustion facilities (25-27,
29,30). Although atmospheric organics vary considerably in structure, many
species are reactive and may be oxidized to form secondary and tertiary
pollutants. One of the most important oxidizing agents in the atmosphere is
ozone produced in photochemical reactions in troposphere and stratosphere
(26). A fraction of the ozone diffuses into the troposphere and contributes
to the background ozone concentration. Oxidation of the organic compounds by
ozone in vapor and liquid phases can result in formation of hazardous and/or
toxic products (2, 3, 25). Under overcast, high-humidity conditions,
dissolved organics may be oxidized in the liquid phase to produce secondary
and tertiary pollutants (26). Rates and mechanisms of the conversion,
however, are not well known. The present research, therefore, was undertaken
to investigate kinetics of oxidation of toluene, naphthalene, phenanthrene,
anthracene, aniline, and a-naphthylamine by ozone in aqueous media.
A stopped-flow sr,ectrophotome ter (Durrum Model D-110) was applied to
conduct the kinetic experiments in distilled water and in aqueous solutions of
various pH values and temperatures. Absorbance data were collected through an
automatic data acquisition system interfaced to the spectrophotometer.
Effects of the temperature and acidity on the oxidation rate were
investigated.
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SECTION 2
CONCLUSIONS
Aromatic amines can he very reactive with ozone as well as reagents of
buffer solutions. In distilled water, aniline reacted with ozone according to
second order kinetics, and the reaction rate constant is nearly independent of
temperature between 10 and 35°C at 2.4 x 10^ 1/M-s. The reaction between
a-naphthylamine and ozone was slightly slower; the second order rate constant
is enhanced by temperature increasing from 0.53 x 10-* to 1.3 x 10 1/M-s in
the temperature range of 5 to 35°C.
The present research confirmed that polycyclic aromatic hydrocarbons in
general are more reactive than simple aromatic hydrocarbons with ozone in the
aqueous phase oxidation. Of the aromatic compounds investigated, the reaction
between anthracene and ozone was the fastest with the second order rate
constant of about 2 x 10 1/M-s in acidic, solutions. The ozonation of
phenanthrene also was fast and the rate constant increased with the pH value
and temperature. At 25°C, the second order rate constant increases from
1.94 x 104 to 4.75 x 104 1/M-s as the pH value varies from 2.2 to 7.0. The
second order reaction between naphthalene and ozone was moderate with the ratt
constant varying from 850 to 3750 1/M-s in the pH range of 3 to 7 at 25°C. As
expected, the rate of ozonation of toluene was slowest among the oxidation
reactions of aromatic hydrocarbons investigated. Similar to the ozonation of
benzene, the order of reaction of toluene changed with acidity indicating a
possible shift in the mechanism of reaction.
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SECTION 3
RECOMMENDATIONS
Kinetic studies of oxidation of aromatic and olefinic compounds in
aqueous media should be continued. The results of this work suggest that many
organic pollutants are very reactive with ozone and that the rate of aqueous
phase reaction can be much faster than the reactions in vapor phase. The
kinetic information is needed in determining reaction life times and r~- - • ..f
degradation of individual aerosols or hazardous and/or toxic pollutants;.
Products of the ozonation, and mechanisms of the oxidation should be
determined. Results of those studies can provide basic knowledge required for
development of control strategies for individual organic pollutants.
Absorption of ozone in water drops and subsequent reactions between
dissolved ozone and contaminants in the liquid phase may play important roles
in detenriining the transport and fate of individual pollutants . Although,
much attention has been focused on the vapor phase reaction, the results of
the present research suggest that the gas-liquid reactions can be very
important in the pollution control because of enhancement of the mass
transport by the aqueous phase reaction. Theoretical and experimental
investigation of the mass transport and chemical reactions in gas-liquid
systems are recommended to provide the vital information.
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SECTION 4
MATERIALS AND EQUIPMENT
PRODUCTION OF OZONE GAS
Ozone gas was produced by passing a stream of extra dry oxygen through a
Welsbach Model T-408 Laboratory ozonator. The czonator is operated on the
corona discharge principle by imposing a high ac voltage across a gap in the
presence of an oxygen-containing gas.
30 > 20 (1)
The efficiency is very low by this method, because only about 10% of the
energy is utilized to produce ozone. The remainder is converted to light,
sound and primary heat and, therefore, heat has to be removed by a water
cooling system. Otherwise, a build-up of high temperature in the discharge
space will reduce the yield of ozone since decomposition of ozone is very
temperature sensitive. The concentration of ozone produced by the ozonator is
affected by several factors, including flow rate of the gas stream, voltage,
gas pressure, impurity of the gas stream and temperature of the cooling system
(18, 21, 23, 28, 33). Water is not allowed in the ozone producing chamber,
and therefore, the use of extra dry oxygen of -60°F dew point is recommended.
The ozonator is operated on 115 volt, 50/60 cycles, single phase power
supply. The oxygen gas pressure has to be regulated to 10-18 psig and
adjusted at 5-8 psig in the ozone producing chamber. The gas stream can be
adjusted and measured at a flow rate between 0.5 to 2.0 1/min. The voltage
can be set between 70 to 115 volts depending on the desired concentration of
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ozone. Tap water is used in the cooling system. To achieve a high
concentration of ozone in the gas stream, a small flow rate, high voltage and
pressure, and low temperature of the cooling system 'should be applied.
The ozone outlet is first connected to an empty washing bottle then to a
second washing bottle containing an appropriate buffer solutions. The
remaining ozone is allowed to pass through a third empty washing bottle then
absorbed in potassium iodide solution contained in a fourth bottle. The use
of Tygon tubing, a good ozone resistant flexible material, is recommended for
all connections in the output streams from the ozonator.
PREPARATION OF BUFFER AND AQUEOUS SOLUTIONS
Buffer solutions were prepared by adding appropriate amount of ACS grade
chemicals such as HC1, NaH2 po4 , Na2HP04 , H-jPO^ or NaOH in distilled water
for control of pH value and ionic strength. Any contaminants in the distilled
water were oxidized by bubbling of ozone gas for at least 30 minutes and the
residual ozone swept by nitrogen gas before preparing the buffer solutions.
Quantities of the various chemicals were determined from the following
Henderson-Hasselbalch equation:
pH= p".a+ log(salt)/(acid) (2)
The ionic strength of a buffer solution is controlled by
I - °-5ZiCiZi (3)
where C^ £s the moiar concentration of component i, and Z^ represents the
number of charges of ion for component i.
In the present research, pH values of the solution were controlled in the
range of 2 to 7 and the ionic strength was maintained at 0.1 in all solutions,
The pH value of a solution was measured by a Sargent-Welch Model NX pH meter.
All buffer solutions were utilized within 2 weeks after preparation.
5
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Highest grade chemicals available froii various suppliers were purchased
for this research. For each compound, a stock solution of a certain
concentration was first prepared by dissolving appropriate amounts in 1000 i 1
of a buffer solution of desired acidity, and the mixed solutions stirred for 5
to 40 hours to achieve homogeneity. The stock solution was then diluted using
buffer solutions of the same pH value to obtain reactant solutions with
concentrations varying from 1.0 x 10 to 1.0 x 10 M. The solutions were
kept in a cool and dark place, and used within a few days to avoid
degradation.
STOPPED-FLOW SPECTROPHOTOMETER SYSTEM
The apparatus useJ in the kinetic experiments is a Durrura-Gibson Model
0-110, stopped-flow spectrophotoraeter. It is a complete system for rapid
mixing of two liquid reactants aid for measurement o£ the change in optical
absorbance as a function of time at a wavelength ranging from 200 to 800 mn.
This system consists of six functional subsystems as showii in Figure 1. The
subsystems are described below;
(1) The sample flow subsystem - includes reservoir syringes, drive syringes,
valve blocks, sample mixing j
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» \\ VAC AUXILIARY
Figure 1- Functional Block Diagram of Stopped-Flow
Spectrophotometer System
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photomultLplier. The log buffer amplifier serves an important function for
the signal display on the oscilloscope or input to the data acquisition
system. A large time constaiv: retting provides a stable signal but a slow
response. On the other hand, a fast response but unstable signal can be
obtained by using a small time constant.
(5) The temperature control subsystem - circulates a coolant (e.g. water)
that maintains a constant temperature for all the parts of the flow subsystem.
In the present research, a constant temperature bath, a Forma Scientific Model
2067 circulating system was used to maintain the circulating liquid at a
desired temperature.
(6) Data recording devices - A Hewlett Packard 1207 A storage oscilloscope, a
Model 198 A oscilloscope camera and a Model 680 strip chart recorder.
The kinetic experiments were conducted under isothermal conditions by
water circulation through the stopped-flow spectrophotometer system to
maintain a desired temperature. The stopped-flow equipment can be operated
following appropriate valve settings and filling of two reactants in the drive
syringes from the two separate reservoir syringes. Then, the valves may be
r°set for measurement of absorbance during a reaction. By activating the flow
acturator,two solutions containing an organic compound and a solution of
dissolved ozone in the two separate drive syringes can be forced to mix
rapidly (99.5% complete within 0.0005 to 0.002 sec.) in the mixing jet and
flow through the observation cuvette. Simultaneously, the light intensity
signal from the photomultiplier tube, transmitted by passing a monochromatic
light through the mixed solution in the cuvette observation chamber, is
amplified and an output signal proportional to the absorbance (10 volts = 1
abscrbance) is generated. The output signal is transmitted as function of
reaction time to the oscilloscope or strip chart recorder for display and to
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an automatic data acquisition system for collection, storage and analysis of.
the kinetic data.
DATA ACQUISITION SYSTEM
A data acquisition system for rapid kinetic experiments has been developed
(15), utilizing the hardware components on loan from U.S. Environmental
Protection Agency. A PDF 8M computer with A/D converter, tape drive unit and
formatter, analog recorder, oscilloscope, and input/output terminal interfaced
to the stopped-flow spectrophotometer provides the complete data acquisition
system. Timing for actual data collection is generated through internal
computer software with inclusion of A/D conversion time. The sampling time
interval can be selected from the range of 1 x 10~^ to 100 sec. depending on
the speed of a reacton. A total of 19-'t3 data points can be collected and
storedj and therefore, the total sampling time for a reaction can vary from
0.1943 to 1.943 x 105 sec.
The system software, along with tape drive programs, is stored in core
memory. All other machine language subroutines as well as high level programs
are stored on tape. User-selected data, special analysis programs, and
results are also stored on t<-.j.e. Typical application of the system begins
with a simple initialization from the computer console. After this start-up,
software programming provides complete interaction between the user and the
DecwriLer III or teletype terminal. The user is directed through the process
by a series of decision steps. The initial step is to choose a data file,
either existing or new. Upon selecting the desired data file, a sequence of
options is -.nitiated, the first one being data collection. If an existing
data file is chosen, the decision to write over this data is left open to the
viser. Should the data collection be negated, the existing data is transferred
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from the tape to core memory for a series of print and analysis options. If
data collection is assumed, various experimental parameters and timing
constants are entered. After each run, the user may then decide whether to
keep the new spectrum or return to re-initialize data collection for the same
data file. A successful experiment may then be catalogued, and results
printed out or stored on tape. At this juncture, the user can choose another
data file or continue the print and analysis options. An IBM personal
computer system also has been interfaced to the data acquisition system. The
kinetic data for any experimental run, therefore, can be transferred to the
IBM computer system for further correlations and analyses. The kinetic data
obtained in this research were preserved and documented i.n several theses (23,
28, 33).
10
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SECTION 5
EXPERIMENTAL PROCEDURES AND ANALYSES
DETERMINATIONS OF ABSORPTION COEFFICIENTS
Preliminary tests were conducted to determine the relationship between the
absorbance of an ozone solutions measured by the spectrophotometer at a given
wavelength and the actual concentration of ozone in the solution. The
iodometric method was employed to determine concentrations of ozone in aqueous
solutions. The oxidation of potassium iodide by ozone results in liberation
of iodine which can be titrated by sodium thiosulfate solutions using starch
as an indicator. The instantaneous reactions in the solution are as follows:
03 + 2KI + H20 -»• 02 + I2 +• KOH. (4)
*2 + 2Na2 S2 fly* Na2S406 t- 2NaI (5)
In a preliminary experiment, the absorbance of a buffer solution
containing dissolved ozone at a fixed pH value and temperature was measured by
the spectrophotoroeter at a given wavelength. Simultaneously, a portion of the
solution was mixed with a solution containing excess potassium iodide. Since
potassium hydroxide was produced, the mixed solution was acidified by sulfuric
acid to maintain a pH value of 2 or below. The mixed solution was stirred
continuously and titrated by the sodium thiosulfate solution until the yellow
color of the liberated iodine nearly disappeared. As the colorimetric end
point was approached, starch indicator was added to develop a dark blue color.
The titration was continued until the blue color disappeared from the
solution. The initial concentration of ozone in the solution, therefore, can
be calculated from the amount of sodium thiosulfate consumed in the titration.
Absorbances of various concentrations of dissolved ozone measured at a
11
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given wavelength indicated that a linear relationship existed between the
absorbance, A, and the ozone concentration, C This result suggests that the
Beer's law is valid within the concentration range of interest.
*=ebCA (6)
In the above equation, the molar absorptivity or absorption coefficient, £ , at
given wavelength with light, path b=2 cm, can be obtained from the slope of the
linear plot. The preliminary test also showed that temperature changes have
negligible effect on the measured absorbances. The experiments were conducted
at various wavelengths yielding the absorption coefficients or molar
absorptivities as follows:
wavelength, nm 245 260 270 280
absorptivity, M"1 cm"1 2,070 3,590 1,650 820
The above results are in good agreement with those reported in the
literature (1, 16, 181.
SELECTIONS OF WAVELENGTHS
Careful selection of optimum wavelengths is an important consideration in
the application of a spectrophotometer. Tn conducting a kinetic experiment,
it is often advisable to select a wavelength at which sbsorbance is exhibited
only by one of the reactants. This will allow applications of simple methods
for analysis of measured absorbances. Also, the error introduced by the uses
of the measured absorbances can be minimized without complications of the
ebsorbance behavior exhibited by other reactants.
In experimental investigations of ozonation of organic compounds in
aqueous solutions, it is desirable to conduct an experiment at a wavelength
12
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where the absorbanca of ozone is predominant in the aqueous solution. By
conducting the experiment in this manner, it is permissible to analyze
absorbances exhibited by ozone and reaction products only and neglect little
absorbances of the organic reactants.
Since strong absorbances are exhibited by ozone in thj aqueous phase in
the wavelength range of 245 to 280 ntn as shown in the previous section,
attempts were nade to conduct the kinetic experiments at a wavelength within
this range. Therefore, preliminary tests were carried out to measure
absorbances of individual organics in the aqueous medium at these wavelengths
to aid in the selection of an optimum wavelength for the kinetic experiments.
As illustrated in Figure 2 for aniline at a given concentration in distilled
water, the absorbance approaches a rainiraum at 260 rim. Thus, the optimum
wavelength of 260 was selected to study che anili\=-ozone reaction. Similar
tests were conducted for other organic reactan.'-s to determine the wavelengths
selecter1 in this project.
SELECTIONS OF TIME CONSTANTS
There is a selector switch in the slopped-f low spectrophotoraeter to choose
one of the six filter capacitors to connect the signal path between the
photovnul tiplie«r tube output and the log buffer amplifier. The output of the
photomul t iplif!r consists of a series of very minute pulses, whose time average
is proportional to the t ransmi ttance of a sample. For a large time constant,
the filter capacitor produces a very smooth curve on the oscilloscope display,
but a rapid change in transmi ttance cannot be followed because of slow
r s jponflo. s . A short time constant permits response to a rapid change ir,
transmi ttancc , but the oscilloscope display may contain a great amount of
13
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CJ
Ci
2.6
2.4
2.2
2.0
1.8
1-6 _
1.2
240
pll = 4.90
Temperature = 25°C
Concentration of Aniline
5 K
H
250
260 270
Wavelength, nm
280
290
Figure 2. Absorbances of Aniline in Water at Various Wavelengths
U
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Each filter capacitor is a first-order element, and its dynamic behavior
is dictated by a first-order differential equation in time with a designated
time constant,T . The output voltage signal of a filter capacitor in response
to a step change in an input voltage, AV^ £s then governed by
AV(t) = AV.(i-exp(-t/T)) (7)
As can be seen from the above equation, the dynamic response of output signal
approaches unity exponentially. To achieve 99.9% response or higher, fo^
example, the time .required from the reaction experiment is t _> 6.9T. This
implies that for a filter capacitor of 10 ras time constant, 690 ms would be
required before achieving more than 99.9% response. If the time constant of a
filter capacitor is I ms, on the other hand, only 6.9 ms is needed to achieve
the same or better response.
Although, a filter capacitor of small time constant is desirable as
discussed above, the response signal may contain certain amount of noises. To
compensate for this adverse effect, therefore, it is recommended to select a
filter capacitor with a time constant about one-tenth of the reaction time
constant. For example with a first order reaction of a rate constant, 10 s ,
the reaction time constant is 1/10 second or 100 ms, and a filter capacitor
with the time constant of 0.1 X 100 or 10 ms should be selected in conducting
the experiment.
Some preliminary experiments for the reaction of toluane and ozone in
neutral solutions were carried out using various filter capacitors to
illustrate the effects discussed above (33). The experiments were conducted
at 25°C, end absorbances of a reaction were followed at the wavelength o^." 280
nm. The absorbance changes with the reaction time are shown in Figure 3 for
15
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System: Toluene-Ozone
pH: 7
Temperature: 25 C
Wavelength: 280 nm
(Toluene}.: 0.0005 M
t, in sec
Figure 3. Absorbance Curves With Different Time Constants
16
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four runs with time constant of the filter capacitors varying from 1 to 50 ms.
Early portions of the absorbance data are plotted in Figure 4 to demonstrate
the effect of the machine time constant on the absorbances measurements. The
figure shows that the measured absorbance using the machine time constants of
1 and 5 ms declined with time, and that the absorbances increased with time
for the filter capacitors of the time constants of 10 and 50 ms. These
results clearly indicated that at the early period of the reaction, varying
degrees of responses to the input signals were exhibited by using the
capacitors of 10 and 50 ins, and that the observed absorbancea should not be
employed to calculate the reaction rate. On the other hand, the output
signals from the capacitors of small time constants (1 and 5 ms) represented
nearly full responses to the input signals even in the early period of the
reaction. Consequently, thesa output signals with the time constants of 1 and
5 ras were trie representation of the absorbances of the reaction solutions,
and all kinetic experiir for che toluene-ozone system were conducted at the
time constant of 1 ms.
KINETIC EXPERIMENTS AND ANALYSES
As described in a previous section, the stopped-flow spectrophotoraeter
system was utilized for tt.e kinetic measurements. An aqueous solution of
ozone was prepared by bubbling a mixture of ozone and oxygen, produced from
the ozonator, into a buffer solution contained in a Corning 3-040 gas washing
bottle. The washing bottle was immersed in an isothermal bath to maintain a
desired temperature. An organic solution was prepared from the buffer
solution of identical pH value according to the procedures discussed earlier.
The two solutions of ozone and organic reactants were stored in separate drive
syringes in the stopped-flow spectrophotoraeter sy^.dm prior to an experiment,
17
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1.2
1. 1
o 1.0
0.5
50 ms
10 tas
;? as
&=6—a=3=z£z=i£;
1 ms
System: Toluene-Ozone
PH: 7
_ Temperature: 25 C
iVavelength: 280 nm
CToluene)Q: 0.0005 M
£7 •""— —
1
0
\0
sec
Figure A. Absorbance Curves During Early Periods Of Reaction
18
-------
and a constant temperature was maintained in the flow section by water
circulation. By activating the flow actuator, the two reactants were mixed
rapidly in the mixing jet and allowed to flow through the observation cuvette
where the mixed solution was stopped abruptly. Absorbances of the mixed
solution (as voltage in the output signals, 10 vdc per absorbance unit) were
sampled at a fixed wavelength in a time interval varying from 1 x 10 to 100
sec. and stored in the PDF 8 M computer memories. The data acquisition system
was applied using regression analyses to correlate the experimental data.
Table 1. Optimum Wavelengths and Time Constants
System
Aniline - Ozone
a --aaphthy lamine - Ozone
Toluene - Ozone
Naphthalene - Ozone
Phenanthrene - Ozone
Anthracene - Ozone
Wavelength, nm. Time Constant, ras
260 5
260 1
280 1
245 1
260 0.1
245 0.1
Optimum wavelengths and time constants were selected in accordance with
the procedures discussed earlier in this chapter. On the basis of the
preliminary test the settings listed in Table 1 were selected to conduct the
kinetic experiments for various systems.
Because of very low solubilities of naphthalene, phenanthrene and
anthracene in water (22,31), the kinetic experiments for these systems were
conducted in the aqueous solutions with ozone in large excess. For all other
19
-------
systems, the organic reactants were present Ln excess in the aqueous media.
In the decomposition reaction of ozone (A) accompanied by an ozonation
reaction between ozone and an organic reactant (B) in the liquid phase, the
overall reactions can be expressed in the following forms:
kl
(8)
k2
A + bB > Qc (9)
The rate of depletion of ozone in the above simultaneous reactions can be
written as
~dcA/dt = kj CA + k2 C™ Cg (10)
where C^ an
-------
circumstance, the rate equation can be rewritten as
-dC /dt=(-l/b)(dC_/dt)
A D
- k2 CAm CBn 01)
If the organic pollutant B is present in large excess in the solution, its
concentration'remains nearly constant during the ozonation reaction. Thus,
i
equation (11) can be approximated by
-dC./dt = k' C,m (12)
where the apparent rate constant is
Integration of equatior. (13) yields,
ln CA/CAO = ~k ' t> for m = 1 (14)
and
(C^/c^)-^1 = 1 -k'(l-m)CAOm"1 t, formal (15)
In the above equations, C^ is taken as tha concentration of ozone at a time
(0.0005 to 0.002 milliseconds) where complete mixing (>99.5%) of the reactants
in the solution is achieved.
To apply the above integration method for analysis of -ibsorbance data
21
-------
obtained from a given experiment, it is necessary to establish a relationship
between the absorbance and concentration of ozone in tht solution. If no
* »
appreciable absorbance is exhibited by any chemical species other than ozone
in the solution, then the concentration of ozone can be computed from equation
(6) using the known value of the absorption coefficient for a given
wavelength. For cases where absorbances of ozone as well as reaction products
are significant and the ozonation reaction is controlled by an overall step
indicated by equation (9) with negligible effect of the decomposition
reaction, the concentration of ozone can be shewn to vary with the absorbance
of the solution as follows (12,18,21):
CA(t)/CAO = (A(t)-A(-)) /(A(o)-A(»)) (16)
where A(°~) is the asymptotic absorbance of the solution measured after
complation of the ozonation reaction. By employing the relationship given in
equation (16), the dimensionless concentration ran be plotted against the
reaction time on a semi-logarithmic or regular scale according to equation
(14) or (15). Regression analyses can be utilized to find a best correlation
of the experimental data yielding the reaction order m and the apparent raf.e
constant, k1. By plotting k1 versus the initial concentration of the organic
reactant, CBQ, on a logarithmic scale, a straight line can be obtained where
the slope gives the order with respect to the organic compound, n. Thus, the
ozonation rate constant, k can be computed frou the in-tercept of equation
(13).
For the phenanthreie -ozone and anthracene-ozone reactions, the
experiments were carried out with ozone in large excess in the solutions.
Equations (17.) and (13) are replaced by
22
-------
-dCR/dt= k'CJJ (17)
and
b CAO (18)
Integrated equations (similar to equations 14 and 15) can be derived, and the
following relationship is valid (10,12):
CB(t)/CBQ = (A(t)-A(«))/(A(o)-A<-)) (19)
Thus, the procedures outlined earlier also can be applied to analyze the
kinetic data.
If the initial concentrations of ozone and the organic reactants are
comparable, it is necessary to obtain a rate expression from direct
integration of equation (11). For the second order reaction (m=n=l), the
integrated equation can be derived as,
In (b+(CBQ - bCAQ)/CA) = In CBQ/CAO + k (CBO - bCAQ)t (20)
Also, it can be shown that equation (16) is valid in relating the
23
-------
diroensionless concentration of ozone, CA/CAO, with the absorbances of the
solution (12), Knowing the stoichioraetric raHo, b, and the initial
concentration of the organic reactant, CBO) t^e tenn {n the left hand side of
the above equation can be plotted against the reaction time, t. The slope of
the straight line plot is k (CBQ _ bCAQ). The absorba.ice data from the
naphthalene-ozone reaction were analyzed by this procedure.
-------
SECTION 6
RESULTS AND DISCUSSION
OZONATION OF AROMATIC AMINES
Aniline and a-naphthylaraine were reactive with buffer reagents such as
H-jPO^, NaHnPO^ and HC1 as indicated by results of the preliminary tests.
Thus, the kinetic experiments were carried out in distilled water (pH values
of 4.9 to 5.2) without buffer to eliminate interferences of the buffer
reagents. The temperature of the reactions was controlled by water
circulation in the range of 5 to 35°C. The absorbance data obtained for the
systems are tabulated in a thesis (28).
Typical absorbance changes during the reaction of aniline and ozone are
illustrated in Figure 5 for three runs. As shown in this figure, the
absorbance of a mixed solution increased very rapidly during the very early
period of 40-60 milliseconds and then declined slowly in the remaining period.
This behavior tends to suggest the formation of intermediate products of high
absorbances in the firs-t period of the reaction and further reaction or
decomposition of tlie intermediates in the remaining life of the reaction.
The kinetic data were analyzed according to the method outlined earlier.
The absorb.inces measured during a reaction can be transformed ir.to the ratio
of ozone concentration at a given reaction time to its initial concentration,
and the regression technique utilized to correlate the data. The results
obtained for the three typical runs were plotted in Figure 6 to demonstrate
that the absorbance data were best fitted by assuming a first order reaction
with respect to the ozone concentration. The apparent rate constants for a /
given temperature, k'( were calculated and plotted against initial
25
-------
1.4
e
o
C
<«
.0
O
en
1.3
1.2
pH value = 4.75-5.02
Temperature = 25°C
Wavelength = 260 nm
Machine time constant
= 5 ms
1.1 Mf- D24038
D24041
D25053
1.0
0 10 20
Reaction time, t x 10J, sec
Figure 5. Absorbance Changes During The Aniline-Ozone Reaction
-------
<
I
D2404! 90.08
D24038 98.69
D25053 118.23
0.1 —
0.03
Reaction time, t x 10 , sec
Figure 6. Determination Of Order With Respect To Ozone
In The Ant]ine-Ozone Reaction
27
-------
concentrations of aniline on a logarithmic scale as illustrated in Figure 7.
The slope and intercept of the straight line were calculated, and the n v^lue
(order with respect to the aniline concentration) at 25°C was about 0.73. For
practical applications, therefore, the ozonation reaction o£ aniline can be
considered as first order with respect to foncentrat ions of both ozone and
aniline, and the average reaction rate constant is estimated to be
2.5 X 105 M~ls-1 at 25°C. Analyses of the kinetic data for the
tt-naphthylaraine-ozone system also indicated that the overall reaction was
second order with the rate constant of 1.3 x 10^ Ms at 75°C.
Average rate constants for all experiments at various temperatures were
computed as listed in Table 2. In the temperature range of 5 to 35°C, the
rate constants vary from 9.4 x 10"* to 2.47 x 10' Ms for the aniline-ozone
reaction, and from 5.2 x 10^ to 1.25 x 105 M^s"1 for the
a-naphthylaraine-ozone reaction. Standard deviations of the correlated data
range Crom 5 to 32%.
28
-------
O.
cx
500
250
100
50
30
20
B
pH value = A.75 - 5.02
Temperature = 25°C
Wavelength = 260 am
1.5
Slope = 0.73
3 9 10
Initial .roncentration of aniline, C x 10*1, M
Bo
Figure 7. Determination 0£ Order With Respect To Aniline
In The Aniline-Ozone Reaction
29
-------
Table 2. Average Rate Constants for Ozonation of Aromatic
Amines in Water
System Temperature, °C
Aniline-Ozone 5.0
10.0
25.0
35.0
o-naphthylamine-ozone 5.0
10.0
25.0
35.0
Rate Constant, 1/M-s
94,000
230,000
247,000
231,000
•2,000
75,000
98,000
125,000
As can be seen from the table, the rate constants for the aniline-ozone
reaction remain nearly constant at about 2.4 x 10-* Ms in the temperature
range of 10 to 35°C. At 5°C, however, the rate constant drops sharply to
9.4 x 10^ Ms. For the a-naphthylamine-ozone reaction, on the other hand,
the rate constant increases from 5.2 x 10 to 1.25 x 10 Ms as the
temperature increases from 5 to 35°C. The Arrhenius equation can be applied
to correlate this temperature effect as shown in Figure 8 yielding an
activation energy of 4.4 kcal/gmole.
Aniline is more reactive toward ozone than many aromatic compounds
because of the highly activating -NH group. Mechanisms governing ozonation
reactions of aromatic compounds have been discussed by many investigators
(4,5,19,20,24,34,35). Similar to phenol, e lectrophilic ozone attack at ortho
30
-------
«*>
o
300
200
c
o
-------
and para positions of the aniline ring raay be possible, though I-anglais, et al
\
(20) suggested that the major attack occurred at the para position. In an
earlier study of ozonation of phenol in aqueous solution, Li (21) found that
the electrophilie attachment at the ortho position was the major mechanism to
produce catechol and o-quinone.
In spite of the above possible difference in the position of
alectrophilic attack by ozone, absorbance behaviors of the two reactions,
phenol-ozone and aniline-ozone, followed a similar trend. For both systems,
the absorbance of a reaction mixture increased rapidly in the initial period
and then declined gradually during the remaining life of the reaction. The
overall kinetics of the ozonation reaction was second order with first order
each in ozone and in the organic reactant. The second order rate constants
obtained for the aniline-ozone reaction are larger than those calculated for
the phenol-ozone reaction. For example, at 25°C, the average rate constant is
2.95 x ICnM s for the ozonation of phenol in buffer solution of a pU value
of 5.2 (21). On the other hand, this work yields the average second order
rate constant of 2.47 x 10^M~ s for the ozonation of aniline in distilled
water {with the pH range of 5.(P.-5.20) at 25°C. This suggests that the
aniline-ozone reaction proceeded at a rate much faster than the phenol-ozone
reaction. Another interesting observation is that the phenol-ozone reaction
was sensitive to temperature changes but the aniline-ozone reaction was nearly
independent of temperature except at very low temperatures.
In the ozonation reaction of a-naphthylamine, the -NH group acts as a
powerful activating group (electron-releasing group)(.34). Evidences (4,14)
suggest that the ozonation reaction of «-naphthylamine is initiated by ozo^.e
attack at the ring containing the NH prour>, find two nolecules of ozone may he
consumed i-i the electrophi1ic reaction to form a diozonide. Further reactions
-------
with ozone or hydrolysis of the diozonide may be expected following the
Initiation step. If the initial attack of ozone is considered to be the major
reaction controlling the depletion ot ozone in a solution, then the overall
reaction between ozone and a-naphthylamine may be considered second order as
confirmed by the experimental results. A comparison of the rate constants for
the aniline-ozone and '"i-naphthylamine-o^one systems reveals, however,
a-naphthylamine is not as r»=ictive as aniline with ozone in the aqueous pha~j.
Possible oxidation of. aniline and a-naphthylaraine by hydrogen peroxide in
distilled water was also investigated. No appreciable changes in the
absorbance of a mixed solution of tht reactants was detected. This result
auggests that neither aniline nor ct-naphthylamine is reactive with hydrogen
peroxide in distilled water without a catalyst.
OZONATIOtf OF TOLUENE
Absorbance changes during the toluene-os.one reactions are shown in Figure
9 for several runs in aqueous solutions of various pH values. Correlations of
the absorbance data by regression analyses revealed that the ozonation
reaction was first, order with respect to the concentration of ozone. As shown
in Figure 10 for the apparent rate constant versus the initial concentration
of toluene, the slopes vary depending upon the acidities of the aqueous
solutions. The order with respect to the concentration of toluene can be
considered unity (n=l) in the solutions of pH values of 2 and 3, one-half
(n=0.5) at the pH value of 5.6, and nearly zero (n=0) in the neutral
solutions. The reaction rate constants were calculated and documented in a
thesis (33), and the average values are surmnarized in Table 3. Standard
deviations of the correlations vary frora 3 to 23%.
33
-------
Run no. p|| (Toluene) o*|0\H (0,)QX10\M
'
10321
2 J1321 3
3 12323 5.6
13322
1.77
1.78
1.6!
Wavelength: 200 nm
Tomporaturo: 25 C
Ionic Strength: 0.1
Figure 9. Absorbance Changes During The Toluene - O^one Reaction
-------
1
0.1
1
o
o
w
0.01
0,005
Cjuse pH Slope
1 2 .918
2 J> .918
J> 5.6 *k~5k
k 7 .181
H S 8 *
a a ^ """
H _^°.a-—ra a a
^^^J* 0
° 0 \^/
..Q /1&2
*S^*/*
^s^ / System: Toluene-
^ A / Ozone
K x \Vavelength: 280 nm
-' % Temperature: 25 C
y^ Ionic Strength: 0.1
- /
f
X
X F 1
1x10"
1x10
-2
CToluene) Q, M
Figure 10. Determination Of Order With Respect To Toluene
35
-------
The effect of the pH value on the average reaction rate constant is
demonstrated in Figure 11 for exper;ments conducted at 25°C. This figure was
prepared by converting the reaction rate constants at different pH values into
the common unit of the second order rate constant. The figure shows that the
reaction rates were relatively slow in strongly acidic solutions. For
example, the second order rate constants vary between 27 to 29 Ms in the
solutions of pH values of 2 and 3, with Vuilf reaction life time of more than
20 seconds as indicated by the absorbanc^ data tabulated elsewhere (33). The
rate constant increases rapidly with the pH value, and in the neutrsl
solutions, the second order rate constants are in the range of 97 to 228
Ms with half lives of less than 3 seconds. Therefore, the reaction in the
neutral solutions was at least 3 times faster than the reaction in the
strongly acidic solutions.
The Arrhenius equation can i>e employed to correlate the dependence of the
rate constant on temperature as demonstrated in Figure 12 and 13. The figures
illustrate that the influence of the temperature on the r.-jte of ozonation of
toluene Is very significant, especially in the neutral solutions. The
experimental results show that in the acidic solutions with pH values of 2 and
3, the reaction rate increases about twice for a temperature increase of 10°C.
In the neutral solutions, however, the rate increases about 10 tiim;s for the
lame temperature change of 10°C.
36
-------
200
150
r\J
w
c
o
o
0)
CD
T3
S-t
O
-a
c
o
o
-------
Table 3. Average Rate Constants For The Toluene-Ozone Reaction
pH Value Temp,°C
2 10
2 25
2 35
3 25
3 35
5.6 25
7 10
7 25
7 35
Order With
Respect
To 0
1
1
1
1
1
1
1
1
1
Order With
Respect
To Toluene
1
1
1
1
1
1/2
0
0
0
Rate Constant, k
10.54 M"1 s~l
29.25 M~ s"1
53.30 M'1 a~l
27.21 M'1 a'1
53.18 M'1 s'1
1.697 M-1/2 s~l
.01972 s"1
.1634 a'1
1.614 s'1
The activation energies can be calculated from the slopes of the straight
lines in Figures 12 and 13. At the pH values of 2 and 3, the activation
energy was found to be 11.22 kcal/g nole with a frequency factor of 4.91 X 109
M s . In the neutral solutions, the calculations yielded nn activation
energy of 31.26 kcal/g rool.i and a frequency factor of 1.79 X 1022 s"1. The
significant difference in the activation energies at various pll values suggest
that the mechanism of ozonation of toluene might be different depending upon
the acidity of a solution.
The ozonation rate of toluene was faster in the acidic solutions and
slower in the neutral solutions than the benzene-ozone reaction (16.17).
Nontheless, the changes in the reaction orders and the possible shift in the
38
-------
100
rt
c
o
o
Q)
4->
rt
OS
20
10
O : pH 2
X: pH 3
3.25
k2
E
f
= fc(-r/RT)
= 1 "i .22 Kcal/g raole
3.35
B x 10
3
o-l
3.55
Figure 12. Effect Of Temperature On The Toluene - Ozone
Reaction In Acidic Solutions
39
-------
Figure 13. Effect Of Temperature On The Toluene - Ozone
Reaction In Neutral Solutions
.0
1.0
0.5
a 0.1
c
o
o
0.05
0.015
\
3.25
= fe
(-E/UT)
E =31.26 Kcal/g mole
f = 1.79 x 1022 s~}
3.35
TT; x 10
3
-------
reaction mechanisms with the acidity were observed in Che ozonation reactions
of both benzene and toluene. In an investigation of ozonation of polyalkyl
benzenes in organic solvents by Nakagawa, et al. (24), the rate constants were
calculated assuming second order kinetics (first order each in ozone and in
aromatic hydrocarbon). The ozonation reaction of toluene was very slow with
the second order rate constant of 0.166 M s in carbon tetrachloride and
0.091 M"* 3 in acetic acid. Also, as reported in the liter.uturfi (25), the
ozonation rate constant in the vapor phase at 25°C is about 7 M s . The
results of the present research indicated, however, the ozonation reactions in
the aqueous media were faster than those in other environments reported by the
previous investigators. The second order rate constants in the ctrongly
acidic solutions (with pH ranging from 2 to 3) at 25°C ara about 27 to 30
M s . Tne rate constant increases as the acidity of tlv. solution decreases,
and therefore, the ozonation rate of toluene in the aqueous r,edia is much
faster than the rate of reaction in carbon tetrachloride and acetic acid by
several order of magnitude. A rate constant of 14 4s for ozonation of
toluene in water at 20°C was reported in the literature (29). This value
agrees well with the rate constants determined in the present work. A
comparison of the kinetic data also suggests that toluene is more reactive
with ozone in the aqueous phase than in the vapor phase.
OZONATION' OF NAPHTHALENE
It has been suggested in the literature (14,34,35) that two moles of
ozone are required to react with one mole of naphthalene to form a diozonide.
Bailey (4,5) confirmed the attack of the ozones on 1,2- and 3,4- bonds of
naphthalene (the honds with the lowest bond-localization energy) from a
product analysis of ozonolysis of naphthalene in methanol. Since the
41
-------
intermediate steps for e leetrophiIie addition occur very rapidly, it is likely
that the overall reaction is controlled by the initial attack of the ozuue
molecules. Therefore, the overall reaction can be considered oecond order
with first order each in ozone and naphthalene concentrations. Analyses of
the absorbance data obtained in the present uor'ic (33) tend to support the
second order kinetics for the reaction between ozone and naphthalene in the
aqueous phase.
Sinco the initial concentrations of naphthalene and ozone were in the same
order cf magnitudes, equations (16) and (20) were applied to correlate the
kinetic data using the stoichioaki-tric ratio of 2. Average second order rate
constants were summarized in Table 4. Standard deviations range from 10 to
272 in these correlations. tt should be noted that the stoichiowtric ratio
of 3 also was suggested by some ?..irlier invewt igators (34,35). Calculated
results (33) indicate that the differences in the rate constant* are
insignificant for the stoichioraetric ratios of 2 and 3.
The experimental results revealed that the acidity of a solution was an
important factor affecting the reaction rate. The reaction in the neutral
solutions was faster than that in the acidic solutions as can be seen iron
Table 4. Temperature is another factor influencing the ozonation rate. For
example, the rate constant increases from 848 to 9845 1/M-s or about 10 times
for a temperature increase of 10° from 25 to 35CC in the acidic solutions
with the pH value of 3.
-------
TABLE 4. AVERAGE RATE CONSTANTS FOR THE NAPHTHALENE-OZONE REACTION
pH value
3
3
5,6
7
_
TEMP,°C
25
35
25
25
RATE CONSTANT, 1/M-s
840
9845
3202
3749
The simplest molecules of polycyclic aromatic hydrocarbons, naphthalene
appears to be more reactive than simple aromatic hydrocarbons such as benzene
and toluene with ozone in the aqueous phasi. For the ozonation of toluene in
the neutral solutions at 25°C, for example, the average second order rate
constant can be calculated to be about 173 1/M-s. This value is much lower
than the rate constant of 3749 1/M-s obtained Cor the ozonation of naphthalene
at the same condition. In the acidic solutions with a pH value of 3, the rate
constants for the ozonation of toluene and naphthalene are 27 and 848 1/M-s,
respectively, at 25°C. Again, this shows that the r^te of ozonation of
naphthalene was much faster than that of toluene. This conclusion is in
agreement with the literature report that polycyclic aromatic compounds are,
in general, more reactive than simple arotaatic hydrocarbons.
OZONATION OF PHENANTHRENE
Since, the solubility of phenanthrene in aqueous media is very low at room
temperature (7xlO~6 gmolo/1 or less), the experiments were conducted with
dissolved ozone in large excess in the solutions. Absarbanca behaviors during
the reaction and various methods for analysis of the absorbance data were
discussed in a recent publication (10). The correlations indicated that '_he
43
-------
reaction was first order with respect to concentrations of both phenanthrene
and ozone. As listed in Table 5, the iverage rate constants range between
14,600 and 101,300 1/M-s indicating that the phenanthrene was reactive with
ozone in aqueous media. The half life of reaction varied from about 0.1 to 1
seconds in the experiments. At 25°C with an initial ozone concentration of
lxlO~^M, for example, the half reaction life decreased from 0.36 Co 0.15
seconds as the pH value increased from 2.2 to 7.0.
The effect c£ acidity on the reaction rate constant is indicated in Figure
14. The rate constant increases with the pH value but the rate of increase is
much more pronounced at 35°C than at 15°C. At 35°C, the rate constant varies
from about 20,000 to 101,300 1/M-s in the pH range of 2 to 7. On the other
hand, the rate constant increases from about 10,000 to 26,000 1/M-s at 15°C in
the same pH range. It should be noted that rate constants reported in a recent
paper (9) for two pH values at 25°C agree fairly well with our results. The
observation in the publication regarding insignificant influence of acidity on
the reaction rate, however, is not supportable frora the findings of the
present work.
The Arrhenius equation can be applied to correlate the dependence of the
rate constant on temperature as shown in Figure 15. Activation energies of
the reaction are estimated to be about 7 kcal/mole at the pH value of 3, and
12 kcal/mole at pH values of 5.8 and 7. The calculated activation energies
and rate constants suggest that phenanthrene was more reactive than some
simple aromatic compounds with ozone in aqueous media. The slight change in
the activation energy with the acidity might be indicative of so;ne differences
in the reaction mechanisms between the phenanthrene and ozone at various
acidities. Tha differences in absorbance behavior in acidic and neutral
solutions also were noted earlier. However, judging from the fact that the
AA
-------
Figure 14. Effect Of Acidity On Tho Phcnantbrene - Ozone Reaction
-------
10
o
3 8
O
E
cn
C
o
CJ
a)
u
rt
Pi
C
o
O
fl3
CD
Pi
N)
Symbol pH Value
Activation
Energy ,
kcal/mole
3.2 3.3 3.4 3.5
Temperature"1 , (l/T)x!03, °K~1
3.6
Figure 15. Effect Of Temperature On The Phenarthrene - Ozone
Reaction
-------
overall kinetics of the reaction can be treated as second order In all aqueous
solutions, the differences in the reaction mechaaisras at various acidities
i
might be of a minor nature. Therefore, the conclusion regarding the major
reactions of phenanthrene by ozo.iolysis at the 9,10-bond by the previous
investigations (8,27,32) appears, reasonable. The results of the above
investigation were discussed in details in the recent publication(10).
Table 5. Average Rate Constants for the Phenanthrene-Ozone
Reaction
pH value
2.21
3.00
3.00
3.00
5.75
5.75
5.75
5.75
7.0
7.0
7.0
Temperature . °C
25
15
25
35
10
; 15
25
35
15
25
35
Rate Constant, 1 /M~=
19,371
14,578
23,672
31,515
15,676
22,312
40,820
86,232
26,157
47,468
101,332
OZOSAT10N OF ANTHRACENE
Similar to phenanthrene, the solubility of anthracene in water is very low
(22,31), and the kinetic experiments were carried out in aqueous solutions
47
-------
with ozone in large excess. Regression analysis confirmed that the overall
reaction was second order with first order each in ozone and anthracene
concentrations (23). Average rate constants for the reaction are listed in
Table 6.
In the aqueous solution with the average pH value of 2.2, the second order
rate constants increase slightly from 1.87 x 107 to 2.68 x 107 1/M-s as the
temperature increases from 6 to 35°C. The Arrhenius equation can be employed
to correlate the dependence of the rate constant on the temperature as
demonstrated in Figure 16. The activation energy can be calculated from the
slope of the straight line in the figure to be 2.5 Kcal/gmole and the
frequency factor is estimated to be 1.72 x 10 1/M-s. The influence of
temperature on the rate of the ozonation of anthracene seems to be unimportant
in comparing with the ozonation race of many other aromatic hydrocarbons. The
rate constants for the experiments conducted in both the buffcjr solutions and
distilled water of different pH values at the same temperature of 25°C are
nearly identical. The deviation among the average rate constants is less than,
10%. These results tend to indicate that the acidity of the aqueous media is
net i; significant factor in controlling the ozonation rate because of the
highly reactive nature of anthracene by ozone attack. This same phenomenon
was observed in studying the very fajt reaction of cyclohexcne with ozone in
aqueous solutions as discussed in a previous work (15,16). Although th~
molecular weight of anthracene is identical to that of phenanthrene, their
rates of ozonation were quite different.
The half reaction life of anthracene varied from 0.13 to l.l mi Ili-secouds
in the experiments. On the other hand, the half life of phenanthrene was
about 0.1 to 1.0 seconds as reported earlier. The reaction of anthracene with
ozone was much faster than the phenanthrene-ozone reaction in the aqueouji
48
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7.0
6.0
5.0
o
f-4
K
« 3.0
c
o
u
c
o
V
a-.
1.5
3.22
pH V«lu« - (2.19 - 2.24)
k2 = f exp ( -E/RT )
E = 2518 cal/gnolc or 2.5 KCal/goole
f -1.72 x 109 M^Sec"1
3.30
3.40
x 10
3.50
3.60
Figure 16. Effect Of Temperature On The Anthracene -
Ozone Reaction
-------
media. This might be attributable to the difference in molecular structures
of the two compounds though the major attack of ozor.e seeowd to occur at 9, 10
carbon positions of both phenanthrene and anthracene. An electrophilic
reagent, ozone attacks carbon atoms 9 and 10 of anthracene which are the
positions not only of lowest atom-localization energy, but also of lowest
para-localization energy. Three moles of ozone were consumed per taole of
anthracene to yield nnthraquinoae (6,7,11,27). In the case of phenanthrene,
the ozone attack resulted in the cleavage of phenanthrene - like double bonds
at 9,lO-positiona. This was a bond of lowest bond - localization energy.
This, resulted in the formation of discidg, and only one mole of ozone was
required for each mole of phenanthrene(8,2?).
The I,2-bond of anthracene represents the bond of lowest bond -
localization energy. Instead of this bond, the major ozone attack occurred at
the 9,10-atom positiona. As indicated by the results of experiments
conducted in organic solvents (6,7), the attack at the position of lowest
atora-localization and para-localization energy was much more efficient than
the attack at the position of lowest bond-localization energy. The
experimental evidences from this work suggested that the rate of ozonation of
anthracene was much faster than the rate of ozonation of phenanthrene. For
example, the rate constant of 2.78 x 10' 1/K-s for the anthracene-ozone
reaction in distilled water at 25°C is about 700 times larger than the rate
constant of 4.1 x 10* 1/M-s for the plienanthrene-ozone reaction at the similar
condition. Even though both anthracene and phenanthrene have very low
solubilities in water, their absorbance behaviors followed different trends.
For the ozonation reaction of anthracene, the absorbance of a mixed solution
increased rapidly in the initial period then declined very slowly during the
remaining life of the reaction. However, in the case of phenanthrene-ozone
50
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reaction, the absorbance declined sharply in the early period with reaction
time and approached an asymptotic value as discussed earlier.
The second order rate constants for the reactions of pyrene, phenanthrene
and benzo (a) pyrene with ozone in water at 25°C are about 4.0 x 10^,
1.5 x 10^ and 0.6 x 10^ l/M.-s, respectively, over the pH range of 1 to 7 (9).
The reaction rate constants for ozonation of benzene, toluene, naphthalene,
phenanthrene and anthracene in the aqueous media have been obtained in this
and previous studies. Among the above organic compounds, anthracene was found
to be the most reactive with ozone in the aqueous phase. The reactivity trend
for the aromatic hydrocarbons investigated in this and previous projects can
be summarized in the following manner:
Anthracene > Phervanthrene > Naphthalene > Toluene > Benzene
TABLE 6. Average of Rate Constants for the Antliracenc-Ozone Reaction
pH Value Temp, C Rate Constant, l/M-s
2.19-2.24 6.0 1,87
25 2.57
5.10 25 2.78
distilled water 35 2.68
51
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