EPA-600/3-76-089
August 1976
REACTION KINETICS OF OZONE WITH SULFUR COMPOUNDS
by
Ronald E. Erickson and Lei and M. Yates
Chemistry Department
University of Montana
Missoula, Montana 59801
Srant Number R-800655
Joseph 0. Bufalini
Gas Kinetics and Photochemistry Branch
Environmental Sciences Research Laboratory
Research Triangle Park, North Carolina 27711
U.S. ENVIRONMENTAL PROTECTION AGENCY
OFFICE OF RESEARCH AND DEVELOPMENT
ENVIRONMENTAL SCIENCES RESEARCH LABORATORY
RESEARCH TRIANGLE PARK, NORTH CAROLINA 27711
-------
DISCLAIMER
This report has been reviewed by the Environmental Sciences Research
Laboratory, U.S. Environmental Protection Agency, and approved for publication.
Approval does not signify that the contents necessarily reflect the views
and policies of the U.S. Environmental Protection Agency, nor does mention
of trade names or commercial products constitute endorsement or recommendation
for use.
ii
-------
ABSTRACT
This report presents data on the reaction between ozone and several
sulfur compounds which are air pollutants. The compounds of interest are
sulfur dioxide, dimethyl sulfide, methanethiol and dimethyl disulfide.
The rate of the reaction between ozone and dissolved sulfur dioxide
is strongly pH dependent. This is so because sulfite ion reacts extremely
rapidly (second order rate constant four orders of magnitude higher than
that of bi-sulfite ion). These results suggest that under some conditions
atmospheric oxidation of sulfur dioxide may involve ozone.
Dimethyl sulfide was found to reaction extremely rapidly with ozone
in the gas phase but reproducible kinetic data were not obtained.
Stoichiometric and yield data from the reaction between ozone and
methanethiol or dimethyl disulfide in aqueous solution indicate those
reactions to be complex mechanistically, although one product, methane
sulfonic acid is predominant.
iii
-------
CONTENTS
Abstract 111
Tables vi
Acknowledgments vii
I Conclusions 1
II Recommendations 2
III Introduction 3
IV Experimental Methods 14
V Discussion and Results 30
References 58
-------
TABLES
Number Page
I Gaseous Sulfur Loss Balance Sheet for Southern Kraft
Mill A 5-6
II Products of Ozonolysis of Dlsulfides 13
III Sulfur Solution Preparation 15
IV Buffer Solutions 16
V Mean Rate Constants 31
VI Equations for Calculations of Specific Rate Constants 35-36
VII Specific Rate Constants 37
VIII Rate Constants Ozone Bisulfite (Penkett) 41
IX Ozonation of Thiols:Stoichiometry 47
X Ozonation of Disulfides:Sto1ch1ometry 53
vi
-------
ACKNOWLEDGEMENTS
Much of the experimental work 1n this report was accomplished by
three graduate students, Robert L. Clark, Robert Moody, and Lawrence
Schmidt, of the University of Montana's chemistry department. Technical
assistance was also afforded by David McEwen.
vi i
-------
SECTION I
CONCLUSIONS
The first significant scientific finding of the research
carried out under this grant is that the rate of ozonation
of sulfur IV in aqueous solution is controlled, almost
exclusively, by the concentration of sulfite ion. Sulfite
reacts about four orders of magnitude faster than bisulfite
ion and at least two million times faster than aqueous sulfur
dioxide.
However it is the absolute magnitude of the specific
rate constant for the ozone-sulfite ion reaction which is of
significance to the possible role which ozone might play in
the atmospheric oxidation of sulfur dioxide. Our calculations
show that under certain specified conditions, the most crucial
of which is the presence of liquid water, ozone may play an
important role in the atmospheric phase of the biogeochemical
cycling of sulfur.
Methane thiol and dimethyl disulfide also react extremely
rapidly with ozone in aqueous solution to yield methane sul-
fonic acid along with several minor products. The reaction is
more complicated, in a mechanistic sense, than the sulfur IV-
ozone reaction which is a simple second order reaction.
Gas phase kinetic studies for the reaction between ozone
and dimethyl sulfide show the reaction to be extremely rapid
and complex (mixed order) but too few data were obtained for
meaningful comment.
-------
SECTION II
RECOMMENDATIONS
The rate constants for the reaction between ozone and sul
fur IV in solution which were obtained in this study should
be considered in any model for the atmospheric oxidation of
sulfur dioxide.
Since ozone is highly reactive with many subtances,
including all negative ions which are oxidizable, and many
positive ions in their lower oxidation states, it would
be necessary to know both the concentrations of such ions
in atmospheric water and their rates of reaction with ozone
in order to establish a reasonable model. Such data are not
available.
-------
SECTION III
INTRODUCTION
GENERAL
Our original goal was to measure the rates of the
reaction of ozone with sulfur dioxide and several malodorous
sulfur compounds (mercaptans, sulfides, disulfides and
hydrogen sulfide). Determinations were to be made in the
gas phase for dimethyl sulfide and methyl mercaptan and in
water for sulfur dioxide, methyl mercaptan and hydrogen
sulfide. Several serious experimental problems and a shortening
of the project period (from three to two funded years)
resulted in a more modest set of achievements. We report
here on a complete kinetic study of sulfur dioxide in aqueous
solutions, on stoichiometric and product data of dimethyl
mercaptan with ozone in water, and on the gas phase reaction
of dimethyl sulfide with ozone.
SULFUR DIOXIDE-OZONE
The importance of sulfur dioxide in air pollution
episodes is well documented. Recent reviews have covered
1234
the subject so well ' ' that only a few references which
are directly pertinent to the possible reaction of sulfur
dioxide with ozone will be discussed here.
The reaction of ozone with sulfur dioxide has sparked
some debate. Gregor and Martin's finding of sulfuric acid
in the oxidation of hydrogen sulfide (presumably through
the further oxidation of sulfur dioxide) was not confirmed
by Cadle when both reactants were in low concentration in
the gas phase . Pitts recommended in the "Technical Report
on Air Quality Criteria for Photochemical Oxidants" by the
State of California Department of Public Health that the
ozone-sulfur dioxide reaction be studied in the presence of
water . Cadle suggested that such a study would be fruitless
-------
but did comment that a study of the ozonation of sulfurous acid
o
in fog droplets might be appropriate .
One very interesting study of the reaction of ozone with
g
sulfur dioxide was carried out by Espenson and Taube . Equation
1 describes the general stoi chiometry of the reaction in water
while the most intriguing discovery was that the sulfate
formed contains up to two oxygen atoms from ozone.
H?0
equation (1) S0 + 0 * HSO + H + 0
2 3 2
While the work described in this report was in progress,
two important papers were published on possible interactions
between ozone, sulfur dioxide, and moisture. In the first,
Cox and Pankett observed that both sulfur dioxide and ozone
decay on the walls of containers, the rate increasing rapidly
te
11
with increasing humidity . In the second Pankett reported
the rate of reaction of bisulfite ion with ozone at pH 5
Since his work is directly comparable to ours, we will reserve
discussion of that paper until later in this report.
DIMETHYL SULFIDE-OZONE
Dimethyl sulfide is a malodorous, water insoluble sulfide
whose ozone oxidation products, dimethyl sulfoxide and dimethyl
sulfone are relatively odorless and water soluble. Although the su
fide is not normally found in appreciable amounts in polluted
air, its presence is definitely noticeable where Kraft pulp
mills operate. Akamatsu's studies show that the recovery
process losses plus blow gas emissions constitute the major
sources although not the only source of a series of mal-
1 2
odorous sulfur compounds. Some idea of kinds of compounds
found, the variability of concentrations, and the extent
of the problem for dimethyl sulfide can be found in Table 1,
taken from Akamatsu.
-------
4
C
V
r
e c
0 «
*W •£
dP f
C
» C
H C
(0 4
i t X
4J ^
orf
o
0)
§
rrt
W
n)
tn
«w
O
o
CJ
o
0
en
D
)
•i
3
)
9
4
J
J
» .
•t
5
> ;
W"
a:
tf !
K
rrt
8:
en
w :
CM
O
0 OOOH 0 OO 0
o H cs m o vo H r- o
• •••• • •• •
m H *M *M m o H CM o
00 O
H
H nvoon o> CMCO vo
p«. CM vo r* in H moo CM
in OOOH o oo o
CM en
^a1 Hinvo^j* H Hr— i o o
o o o T* ** o oo o r^
o oooo o oo H n
H VOVOC4T H I^H VD O
O O H M-l O
H
vo oOrf o oo vo u ro
n) CM
0)
C -H
o 0) to 19 ,.^ H^CW 4JC-H JJU-I
>-(> -HQJflJrt («-4,X>i^4-)4J^4')V!'O>M (QU-IP
>OC 0) £OC 4JOIO
OM(dOJ «J
p4 CJ WJ CD OQ PI fjQ CQ
04
•8
•O
-rl
r
I
ca
-------
C
<
1
e ;
o <
M I
ii t
*W
r<
OP t
«
(
%
rH f
id c
-P 4
0"-
b _f
t^ »•
f
C
iH
to
0)
•_J
i
CO
id
rn
".
«M
O
e
ti
w
o
fa
•'
0)
u
M
2
o
(0
1) '
j ;
•4
3 •
5
Q ;
^ •
3
a •
u ;
•4 '
3
j
5-
•4
•i
K
CO
rn
3
«
w
a
2
r*|
W
CN
jTl
M*
CN
O
(0
o ooooo o o
O» OCMrHfNin m tN
rH O in 00 (N O C4 00
^P ro .
vo rHtfCNinvo a\ n
O4 rHr-OOO CN O
in OOiH^fO O »H
in omino rH o
rH OOrHOO O O
O O O O CN O ' O O
O CNOCOOCN O O
O O O CO T* O O O
O OOOrHO O 0
o r»vomoiH iH CN
m ocNOino o o
o ooooo o o
•H CNTTrHon r- iH
M) OCNOrHO CN O
m ooooo o iH
.0 o in |. i i i. i
CN 0 rH « 1 1 I I
• •ill 1 1
rH 0 0 1 • 1. 1 | I
O* V< I -H
C O -H rH
>i-H ^ -P rH M-i C (fl fi
>-l> 0) id -rH -HOMO
OrHj^ 4J ^O'O'H-P'H
> O C 0) O -H -P 3 -P
O W (d C Q) ^ (XrH O id 0) id
OW4JrH frOOlrHriJO^N
0)-rt ^j.HrH>ld
« Q U5 Q « pq E-«
0
o
o
o
•H
M3
in
•
CM
rH
in o
^r in
CN O\
iH
r* o
CN . rH
CN 00
rH
o o
CN VO
• *
rH en
• •
rH 0
rH
rH
id
4J
OX «W
•P O O \*
rH Id 3
id «w 0) B -H
-------
The first reported ozonation of organic sulfides was in
13
1933, when Maneck ozonized benzyl sulfide in CC1. and
reported intltlal reaction products to be the sulfoxide
and sulfone. Prolonged ozonation oxidized these initial
products to sulfuric acid, carbon dioxide, and benzoic acid.
In 1942, Bohme and Fischer ozonized several organic
sulfides, including dimethyl sulfide, in chloroform. A
nearly quantitative yield of the sulfone was obtained for
the compounds Me2S, Et2$, (ClCH2CH2)2S, MeS0, EtSCH20,
and (0CH2)2S. That the sulfoxide was formed as a precursor
to the sulfone was shown by the isolation of benzyl sulfoxide
from the sulfide when an insufficient amount of ozone was
used. For some sulfides, i.e., chloromethyl ethyl sulfide,
the sulfoxide was the final product.
While studying the possibility of the use of ozone as
a titrimetric agent for quantitative determination of
1 5
olefinic unsaturation in petroleum fractions, Boer and Kooyman
ozonized several sulfides to determine the effect of their
presence in the olefinic mixtures. The dialkyl sulfides were
reported to consume from ^ to - mole of ozone per mole of
i* <•
sulfur compound. Benzothiophene consumed exactly one mole
of ozone, while thiophene and dibenzothiophene did not inter-
fere with the reaction.
In 1954, Bateman and Cunneen studied the autoxidabi-
1ity of several monosul fides at temperatures ranging from 45°-
75°C and found that saturated alkyl and phenyl sulfides
absorb no oxygen during 24 hours at 75°. Similarly inert
were the benzyl and diphenylmethyl substituted sulfides.
While the benzyl sulfides undergo photo-oxidation, the
dialkyl sulfides oxidized only when catalyzed by aa1-
azobis-isobutyronitrile. The diphenyl sulfides were inert
to all autoxidation.
-------
Barnard found that monosulfide-ozone reactions in solution
at -25°C yielded not only the sulfoxide and sulfone but
other products as well. For the reaction:
R1 • S • R" + 203 ->• R" • S02- R" + 202
less than the theoretical amount of ozone was needed for
a nearly quantitative yield of the sulfone. The ozonation,
run just to completion, of di-n-butyl sulfide gave a 100%
yield of the sulfone "smelling of butyraldehyde and butyric
acid and reacting acid to litmus". The reactivity of the
sulfides indicated that the most easily oxidized sulfides
required the least ozone.
Since saturated alkyl sulfides and several aryl sulfides
are inert to molecular oxygen without a catalyst, Barnard
speculated that the ozone must either catalyze the sulfide-
oxygen oxidation or produce, after loss of an oxygen atom,
activated oxygen molecules some of which participate in
the oxidation of the sulfide or sulfoxide.
Barnard's ozone uptake curves indicated that oxidation
occurs in two stages to give first the sulfoxide and then
the sulfone. By following the ozonolysis of cyclohexyl
methyl sulfide with infrared spectroscopy, he was able to
show that at least 98-99% of the sulfide was converted to
the sulfoxide before any sulfone was detected. This indicates
the rate of sulfoxide formation is 50-100 times faster than
that of the sulfone.
1 8
More recently, Maggiolo and Blair found that ozone
reacted with dialkyl and diaryl sulfides according to the
equation:
° °
-------
When the reaction was run in a nonpolar solvent, it was
found to be stoichiometric for the aliphatic and aromatic
sulfoxides and for the aliphatic sulfones.
.. S .- CH2-0
10. 7g
(.0499)
2.33g
(.0485)
CHC
20°C
m .
11. 3g
(.0491)
0-CH
CH2-
CH2-0
10. 7g
(.0499)
4.71g
(.0981)
12. Og
(.0487)
20
1 9
Horner, Schaefer and Ludwig also found that sulfides
react in ethyl chloride with ozone according to Maggiolo
and Blair's equation to give sulfoxides and sulfones in 80-
100% yield.
Six years later, work by Hughes, McMinn, and Burleson
lent still further support to Maggiolo and Blair's stoichiometry
by obtaining stoichiometric yields of sulfoxide and sulfone
when Bis (B-Hydroxyethyl ) sulfide was ozonized at 100°C.
Research was carried out in this laboratory on the ozona-
tion of dimethyl sulfide in methylene chloride at -78°C.
Nuclear magnetic resonance analysis of aliquots taken during
the reaction revealed dimethyl sulfone formation after only
1% of the reaction was completed. While this would seem to
indicate a DMSO/0- reaction considerably faster than that
-------
for DMS and ozone, it is also possible, due to poor mixing
in the reaction vessel, that pockets of high ozone concentra-
tion could occur and that the sulfone resulted from the slow
oxidation of the DMSO trapped in these areas. Product
analysis indicated an uptake of 1.86 moles of ozone per mole
of sulfide for conversion to the sulfone.
Not until 1966 was there any work done on the vapor
phase ozonation of dimethyl sulfide. In that year, pollution
21
control work resulted in two studies by Akamatsu, et. al.,
on the removal of dimethyl sulfide from pulping processes
by vapor phase ozonation. This work showed that dimethyl
sulfide oxidized with oxygen containing 7% ozone gave
75% dimethyl sulfoxide and 25% dimethyl sulfone. Another
1 p
project of Akamatsu showed that an 85% yield of the sulfoxide
could be obtained when dimethyl sulfide and oxygen containing
20% ozone were fed into a reactor in a 3:7 ratio at 30°C.
22
Hales studied the gas phase kinetics of hydrogen sulfide
with ozone. The results of that study have only an indirect
bearing on the reaction of ozone with dimethyl sulfide, but
the methodology seemed to be particularly useful to the
study we planned to carry out.
THIOLS AND DISULFIDES-OZONE
The first ozone oxidation of a thiol was reported in 1933
I O
by Martin Maneck . He ozonized ethanethiol and obtained
ethanesulfonic acid. The conversion of thiols to sulfonic
acids has also been observed by several other authors.
23
Asinger and coworkers used ozonation of thiol to prepare a
number of long chain aliphatic sulfonic acids in yields from
70 to 100%. Barnard reported quantitative conversion of
thiophenol to benzenesolfonic acid. The conversion of iso-
butanethiol to isobutanesulfonic was accomplished with a 90%
24
yield by Erickson
10
-------
Both Barnard and Erickson have determined the stoichio-
metry involved. Barnard found that 1.6 moles of ozone per
mole of thiophenol were absorbed when the reaction was carried
out by adding ozonized oxygen to a CC1, solution of the thiol
o 24
maintained at -20 C. Working in methanol at -78°C Erickson
found a stoichiometry of 1.6:1 when an ozone-oxygen stream
was used and a stoichiometry of 3.3:1 when an ozone-nitrogen
stream was used for the conversion of isobutanethiol to
the corresponding sulfonic acid.
Disulfides and thiolsulfonates have been reported as minor
13 24
products of the reaction. Maneck found both. In Erickson's
work the disulfide was formed, but the thiolsulfonate was not
identified. However, all the minor products were not identified
Erickson also found methyl isobutanesulfonate, a product
that must have arisen through a reaction with the solvent
I O
(CHoOH). An ozonation of thiophenol performed by Maneck
yielded no sulfonic acid. The only products were the cor-
responding disulfide, thiolsulfonate and disulfone. Maneck
also noted that continued ozonation of either ethanethiol
or thiolophenol resulted in an attack of the carbon-sulfur
bonds and the formation of sulfuric acid.
The mechanism of the ozone-thiol reaction has hardly been
dered. Maneck offe
he gave no proof for them.
considered. Maneck offered the following equations, but
EtSH + 03 -»• EtS03H (22)
2EtSH +0 •* EtSSEt + H90 (23)
0 *
EtSSEt + 20 -> EtSSEt (24)
0
At that time thiolsulfonates were believed to have a disulfoxide
structure and the thiolsulfonate was reported as a disulfoxide.
Two studies of the ozone-thiol reaction in the gas phase
have been carried out. The use of ozone to control Kraft
11
-------
1 2
pulp mill odors has been considered by Akamatsu . He found
that methanethiol was converted to an odorless, water-soluble
compound.
A kinetic study of the ozone-ethanethiol reaction in
the gas phase was performed by Kirchner, Kastenhuber and
25
Biering . Concentrations were in the parts per million
range and under those conditions the reaction proceeded with
carbon-sulfur bond cleavage. The kinetics were determined
by competitive reactions with hexene and isobutene. At
1
30°C a rate constant of 2.4+0.8 X 10 I/mole sec was found.
A series of disulfides has been ozonized by Barnard .
The reactions were carried out in carbon tetrachloride at
-25°C. Sulfonic anhydrides and thiolsulfonates were the major
products. Disulfones were also formed in small quantities.
His results are summarized in Table II. Dimethyl and diben-
zyl disulfide were atypical. Dibenzyl disulfide reacted
by giving carbon-sulfur bond cleavage. Dimethyl disulfide
gave abnormal product ratios. He proposed that the reaction
proceeds according to the following scheme:
DC CD
Ko
J/
RS
°?
3
' ^^
nc.p
~-~-^_
^^.
I
0
RS>1SR
0-0=0
*• DC CD
••• Ko-oK — — — — —
"0-0=0
^ 1/2 RS02SR +
-^Q
o pen c P
^-^^ RSOSOR
0
j fA C.R
0-0-0
R
n c . C .
Ko . O .
o-ozo
+ -
1/2 RSSR
3 \pcn en R
R
R:
Kb
';)
R
\J
I
RS02OS02R
1
12
-------
TABLE II
PRODUCTS OF OZONOLYSIS OF BISULFIDES
17
R 03 Absorbed
(moles)
Ph 2.6
p-Cl C6H4 2.9
n-Bu 2.6
Me 2.0
CH2Ph 5
Sulfonic Thiolsulfonate
Anhydride (%) (%)
90 6
84 10
80 10
39 50
__
Homer et. ajk have ozonized some of the same disulfides as
Barnard. The reactions were done in ethyl chloride and they
obtained different results. Homer obtained an 82% yield
of benzyl phenylmethanethiolsulfonate from dibenzyl disulfide.
Diphenyl disulfide also gave its thiolsulfonate, but only in
23% yield. The product from ozonation of di-n-butyl disulfide
decomposed and was not identified.
13
-------
SECTION IV
EXPERIMENTAL METHODS
OZONE-SULFUR DIOXIDE AQUEOUS PHASE KINETICS
INSTRUMENTAL METHOD
We determined the rate of ozonation of total sulfur
species present in buffered aqueous solutions of sulfur dioxide
at variable pH by a stopped flow technique. The stopped
flow system used was an Aminco-Morrow stopped flow
spectrophotometer with associated ultraviolet source and
a Biomation transient recorder Model 610 with a Cathode
Ray Tube display. The transient recorder storage was then
fed into a Bausch and Lomb VOM-10 recorder. This enabled
the signal collected over a few milliseconds to be recorded
on a 10 second virtual time base for a permanent record,
rather than photographing the CRT display. The log mode (0-1
Absorbance units on the oscilloscope and transient recorder)
was used on all runs. Five traces were recorded with each
run and three runs were made with each pair of reactant
solutions.
TEMPERATURE REGULATION
A large (4 gal.) water bath was used with a cooling coil,
heating element and thermostat.
For the 16°C runs the temperature was maintained with
ordinary cold tap water circulating in the bath. The flasks
containing the reactant solutions were brought to 16° C
before a run was made. Dry air was pushed through the obser-
vation cell area to avoid condensation on the quartz observation
cell.
Temperatures were noted in the bath, before and after
entering the cooling area. The fluctuation was never greater
than 0.2°C between the inlet and the outlet of the cooling block,
14
-------
SULFUR REACTANT SOLUTION PREPARATION
The sulfur solutions were made from sodium bisulfite
(NaHS03) at these pH levels: 0.620, 2.50, 3.55, 3.22, 4.02.
Weighing solid sodium bisulfite was also the procedure used
for all four of the 16°C runs; pH: 0.59, 2.13, 2.09, 3.74.
For the other pH levels (-.301, 1.20, 1.71, 2.11, 2.80,
3.12) S02 gas was bubbled into the buffer solution.
TABLE III SHOWS APPROXIMATE FIGURES OF THE AMOUNT (MLS) NaHS03
ADDED TO A LITER OF BUFFER
TABLE III
pH Molarity NaHS03 Ml NaHS03
0.620 .1 M 5 ml
2.50 .1 M 2 ml
3.55 .01 M 1 ml
4.02 .001 M 5 ml
3.22 .001 M >1 ml
The concentration of sulfur dioxide was determined by
the absorbance of the sulfur solution in the stopped flow
apparatus. Wave lengths and extinction coefficients were
270 nm (e = 461) and 276 nm (e = 500). The concentrations
calculated at the two different wavelengths were always
within 3-4% of each other.
Knowing the sulfur dioxide concentration it was possible
to calculate the concentrations of bisulfite and sulfite ions
by proper manipulation of the known equilibrium constant
15
-------
expressions.
OZONE REACTANT SOLUTION
Ozone solutions were prepared by bubbling 03 into the
proper buffer with nitrogen from ozone saturated silica
gel at -78°. The concentrations were determined at wavelengths
of 250 nm and 276 nm. Extinction coefficients were 2430 and
2150 respectively.
BUFFER SOLUTIONS
In the pH range of 2-5 mixture of 0.10m H3P04 and 0.10m
KH2P04 solutions were used. The actual pH was found with a
Beckman pH meter using the expanded scale. The pH meter
was zeroed in with known solutions in the proper pH ranges.
TABLE IV SHOWS THE PERCENTAGE COMPOSITIONS OF THE BUFFER SOLU-
TIONS
TABLE IV
BUFFER SOLUTIONS
APPROXIMATE PERCENTAGES (0.1M EACH)
% KH2P04
4 2 98
3.5 5 95
3 12 88
2.5 50 50
2 90 10
For pH levels less than 2 0.1-1M H2S04 was used,
16
-------
OZONE-DIMETHYL SULFIDE GAS PHASE KINETICS
INSTRUMENTATION
The stoichiometric data were obtained using a Packard
Model 7300 gas-liquid chromatograph equipped with dual flame
ionization detectors, dual electrometers, and a Vidar 6300
i
Autolab digital integrator. The column employed was a 5' X — "
glass 5% GE XE 60 on Chrom Q. Triple-distilled butyl
benzoate was utilized as an internal standard for determination
of product yields.
Product indentif ication was achieved by comparison of
product g.l.c. retention times with those of known compounds
under identical conditions and by mass spectroscopy using a
Varian Mat III GS/MS system.
APPARATUS
A Welsbach T-408 electric discharge ozonator was used
to produce the ozone. The ozone flow rate was determined
by passing the ozone-oxygen (nitrogen) stream through a
potassium iodide sol ution for a known time and titrating the
solution, on acidification, with a .005 N solution of sodium
thiosul fate.
The equation for the flow rate is:
V = milliliter of NapS-O-, used
N = normality of Na^S^Oo used
t = time in minutes that 0~ was passed into
KI trap
In this study, the carrier gas was simply passed through
the Welsbach ozonator, split to decrease the ozone concentra-
tion, brought to the desired temperature of either 25°C of
17
-------
35° C and Introduced Into the mixing chamber.
The ozone stream was split in order to decrease concen-
trations to those desired. Attempts to decrease ozone
concentrations to those required by decreasing the current
on the ozonator yielded inconsistent ozone flow rates.
After the ozone stream was split, the flow rate of the
resulting stream was measured by a Fischer and Porter
"Tri Flat" rotometer. The nitrogen flow rate was determined
by a Gilmont Model R795 rotometer.
Dimethyl sulfide was introduced into the nitrogen stream
by a .1 ml Hamilton gas-tight syringe drived by a Sage model
341 syringe pump. Flow rates for the DMS were varied from
.0003 to .00006 ml/min.
DMS-03/02 Reactor System
Before the DMS was injected, the nitrogen stream was
brought to the desired temperature in a constant-temperature
water bath. In order to achieve complete DMS vaporization
and to prevent the liquid DMS from splashing on the walls,
the injection chamber was expanded, the diameter of the glass
tubing carrying the nitrogen was decreased and the end was
splayed. After injection, both the N2/DMS and the 03/02
streams entered a styrofoam constant temperature cabinet,
passed through I1 and 17' of glass tubing respectively,
came together in the gas mixing chamber and then entered
into the reactor. The mixing chamber was designed according
22
to Male's specifications and consisted of a Teflon collar
mounted within a stainless steel sleeve. The 03/02 stream
passed into the sleeve and proceeded into the reactor through
four .5 mm radial holes drilled into the collar. The boundaries
between the collar and the sleeve were kept gas-tight using
Vitron "0" rings. The reactor was constructed of precision-
bore Pyrex tubing having an inside diameter of 1/8" + .0002"
18
-------
and had an inside volume of 1.988 mill mters. The reactor
temperature was controlled by means of a water jacket.
Figure 1 diagrams the reactor system. Extreme care was taken
in cleaning the reactor and associated glassware. The pro-
op
cedure followed for cleaning the reactor was that Male's .
The reaction products were collected by means of a freeze-
out trap consisting of a glass spiral connected to an adapted
10 ml pear-shaped two neck flask. Products were then removed
from the trap by warming to room temperature and washing
three times with three milliliter portions of methylene chloride.
A 1 ml volumetric flask was attached to the bottom of the
flask in order to obtain accurate 1 ml volumes of CI^Clp-
product solutions on evaporation. The trap was immersed
in a Dewar containing a solid-liquid slush of n-pentane.
Two other solid-liquid slushes were tried; however, tests,
proved that product loss resulted when either the dry ice/
acetone or the isooctane/1iquid nitrogen slushes were used.
Since the amount of product was small, the 10 ml of methylene
chloride needed to wash out the trap resulted in product con-
centrations too low to be analyzed. Vapor pressure data for
DMSO and CH2C12 were obtained from Karchmer's The Analytical
Chemi stry of Sulfur and Its Compounds and from the Handbook
27
of Chemistry and^ Physics respectively. Extrapolation of
Karchmer's data yielded a vapor pressure for DMSO of less
than .1 mm of Hg at 0°C. A value of .0896 mm of Hg was obtained
p O
using the values of Jose et. al. Methylene chloride's
vapor pressure at this temperature was approximately 129 mm
of Hg. Since methylene chloride's vapor pressure was almost
1500 times that DMSO's, evaporation without product loss
appeared quite feasible as a means of obtaining product
concentrations which could be analyzed. This evaporation
was carried out by removing the 10 ml flask from the rest of
19
-------
o
<0 4J
o 2 ^
N 0)
° g
O
>1
•H
04
O.
3
CO
a
<
f
[
W 4J
K
^oOOOOiT'
j
[^
j
v
f
J
c
^
I
J
<0 n)
Constant Temperature Freez
Cabinet TC
20
-------
the CHpClp-washed trap and attaching a condenser and a
and a nitrogen source. This was then placed in an ice bath
and nitrogen was passed over the solution until a volume
of 1 ml was reached.
In order to determine the ozone uptake of the reaction,
ozone concentrations were measured before and during the
reaction by means of a KI sampler attached to the freeze-
out trap. It was found that a normal glass frit would hold
varying amounts of I2, causing inaccuracy in the determination
of the Oo concentration. It was not possible to divert the
ozone stream in order to wash the frit, since this caused
a pressure change and a resulting fluctuation in the ozone
flow rate. Studies indicated that it took as long as an
hour to restablish the prior CU flow. This problem was
solved by exchanging the frit with a glass tip consisting
of a tapered end and four side openings for a 5-way dispersion
of the gas.
Trapping all of the DMS in the freeze-out trap proved
impossible. Therefore, the DMS entered the KI solution when
the ozone stream was analyzed. DMS reacts with I9 forming
26
DMSO and DMSOp thereby consuming the !« used for ozone con-
centration determination. Tests determined that adding starch
solution to the KI
2KI + 03 + H20 -»• I2 + KOH
CH3SCH3 + I2 + CH3SCH3
CH3SCH3 -I2 + H20 -»• CH3SCH3 + 2HI
0 slow
sample while analyzing for ozone eliminated this problem
since the starch complexed immediately with the I2 on its
21
-------
formation, thus preventing the reaction with dimethyl sulfide.
Water was placed in the sampler between KI analysis in order
to maintain a constant pressure head.
REAGENTS
Nitrogen
Before entering the reaction system, the nitrogen was
further dried and cleaned by passing it through a column
packed with a layer of Drierite and silica gel and a layer
o
of 4 A molecular sieves.
Oxygen
Tests were run on the purity of the oxygen by passing
the gas through a Drierite-si1ica gel column. Extended use
indicated a level of purity and dryness sufficient for this
investigation.
Dimethyl Sulfide
The Baker grade sulfide was dried over anhydrous calcium
sulfate and triple distilled using a 600 mm Vigreux distilling
column attached to a 400 mm Claisen-Vigreaux distilling head.
Dimethyl Sulfoxide
The Baker analysed DMSO was stirred at 25°C with 5% of
its weight of Darco G-60, filtered, treated with MgC03 to
remove the acidic impurities, stored overnight over Cal^j and
triple distilled through a 400 mm Claisen-Vigreaus column
at approximately 10 mm of Hg.
Methylene Chloride
The Mallinckrodt Spectr AR grade methylene chloride was
washed with concentrated sulfuric acid followed by dilute
sodium hydroxide and finally water. The washed material
was allowed to stand overnight over sodium hydroxide pellets
and calcium chloride and was then triple distilled using a
600 mm Vigreux column attached to a 400 mm Claisen-Vigreux
22
-------
distilling head.
Butyl Benzoate
The Matheson Coleman and Bell butyl benzoate was triple
distilled using a 400 mm Vigreux column.
All chemicals were analyzed by gas chromatography for
purity and stored under nitrogen until needed.
Cal ibration o_f Equipment
Calibration of several components of the experimental
system was required before runs were made. These calibrations
are as follows.
Rotometers
Both rotometers were calibrated using soap-film flow
meters. Hales found that on continued use the sapphire
rotometer floats picked up static electricity rendering them
useless for accurate flow indicators. Therefore, the rotometers
were used only as secondary indicators and soap-film flow
meters were used for actual gas flow rates.
Syringe Drive
The syringe pump output was calibrated for several
settings on ml/hr for a .1 ml Hamilton syringe by measuring
the time required for .01 ml of material to be injected.
Five settings were calibrated for dimethyl sulfide (see
below).
Setting on ml/hr Flow Rate
9 3.2593 x 10~4 ml/min
8 2.131 x 10"4 ml/min
DMS 7 1.359 x TO"4 ml/min
6 9.147 x io"5 ml/min
5 6.222 x IO"5 ml/min
23
-------
Determination of Peak Height vs. Concentration Ratios
Gas chromatograph peak height ratios were calculated for
DMSO, dimethyl sulfone, and butyl benzoate. On a mole basis
these ratios came out to be:
Butyl Benzoate 1.0000
DMSO .1359
DMS02 .1788
Evaporation Trap
A known amount of DMSO was added to 10 ml of methylene
chloride and evaporated in the trap to a volume of 1 ml.
Direct comparison of standard gas chromatographic peak areas
with those from the trap indicated a trap efficiency of 99.0
percent.
0.3/0-2 Flow Rate
Since a normal run lasted approximately three hours,
tests were run to determine the consistency of the ozone flow
over this time period. Due to the slow response of the ozona-
tor, a two hour warm-up period was allowed before samples were
taken. Results indicated an average percent deviation of
22
+2.40 percent. Hales conducted experiments to determine the
extent of ozone decay and found that at 28.5°C the decay
was less that 1% during a run. Since our flow rates were
much higher and the reactor diameter smaller, ozone decay
should present no problems.
OZONE-DIMETHYL DISULFIDE AND METHANE THIOL OZONATIONS
ANALYTICAL METHODS
Ozone
The standard iodide-thiosulfate method was used to deter-
mine ozone concentrations. Although there recently has been
some debate concerning the stoichiometry of the ozone-
24
-------
29-31
iodide reaction, the following is generally accepted.
03 + H20 + 21" •+ 02 + 20H" + I2
The iodine was determined by titration with sodium thiosulfate
as shown.
2S203- + I * 21- + S406-
A 100.0 ml aliquot of an aqueous ozone solution was pipetted
into a potassium iodide solution. The solution was acidified
with 5 ml of 6M HC1 and titrated with 0.005 M sodium
thiosulfate. Just before the end point, 3 ml of a 0.3% starch
solution was added and the titration was continued to the
starch end point.
The potassium iodide used was Baker Analyzed and Mallin-
ckrodt AR; both contained less than 0.0003% iodate. Baker
Analyzed sodium thiosulfate pentahydrate was used. The crys-
tals were used as a primary standard after determining that the
values obtained based on the weight of the crystals agreed
within 1.5% of those based on standardization against potassium
dichromate.
Methanethiol
Solutions of methanethiol were made by passing the thiol
into water and then cooling to the reaction temperature (0°C).
An accurately measured aliquot of the thiol solution was
transferred to a glass stoppered flask containing 25.00 ml
of a known concentration (approximately 0.15 M) iodine solution,
After forty minutes, the solution was acidified with 5 ml
of 6 M HC1 and the remaining iodine was titrated with a
standard 0.1 M thiosulfate solution to a starch end point.
This method is based on the following reaction.
2CH3SH + I2 -»• CH3SSCH3 + 2HI
25
-------
Me thanes u_1 f on ic Acid
Potentlometrlc titratlons (Beckman Century SS pH meter)
using standard base solutions as titrant were used to determine
methanesulfonic acid. The base solutions were standardized
against potassium hydrogen phthalate. Ozone was found to
interfere with the determination and was removed by bubbling
out with nitrogen before the titration.
The methanesulfonic acid was identified by converting
it to methanesulfonyl chloride using thionyl chloride. About
2 ml of the suspected sulfonic acid was isolated from the
reaction mixture by evaporating the water and refluxing
with 20 ml of thionyl chloride for four and a half hours.
The thionyl chloride was distilled and the reaction product
was identified by coinjection gas chromatography and its mass
spectrum.
Sulfuric Acid
The precipitation of barium sulfate from aqueous solutions
is the basis for a number of qualitative and quantitative
techniques for sulfate. Qualitative tests for sulfuric
acid were performed using three techniques. One, a small
portion of the solution to be tested was added to a barium
chloride solution; the formation of a white precipitate
indicates sulfate and the absence of the precipitate indicates
the absence of sulfate. For very dilute solutions the formation
of a precipitate can better be detected using a spectrophotometer
Two, a small portion of the solution to be tested was used to
adjust a Spectronic 20 UV - visible spectrophotometer to
100% transmittance and then a few crystals of barium chloride
were added. A decrease in the transmittance indicates pre-
cipitation. Three, a drop of a saturated potassium perman-
ganate solution and three drops of the solution to be tested
were mixed. One drop of the mixture was placed on a filter
26
-------
paper which had been impregnated with barium chloride and
heated at 70°-80° C for seven to eight minutes. The filter
paper was then washed with water and 1 N oxalic acid. If
sulfate was present, the precipitated barium sulfate would
trap potassium permanganate in its crystal structure. The
trapped potassium permanganate would not be washed away;
33
a pink or purple spot indicates sulfate.
Methanesul f im'c Acid
Ferric ions precipitate sulfinic acids as shown.
3RS02H + FeCl3 + (RS02~)3 Fe*3 + 3HC1
Qualitative tests for sulfinic acids were carried out by
adding 1 ml of 15% FeCl3 to 2 ml of the solution to be tested.
OZONATIONS
Ozone-Gas Streams
An ozone-oxygen stream was passed into a 0°C methanethiol
solution in a gas washing bottle equipped with a glass frit.
It was found that large quantities of the thiol were removed
with the exit gases. Cold traps cooled to -78°C and precip-
itation as its silver mercaptide were used to trap the escaping
thiol. The exit gases from the reaction vessel were passed
through an aqueous potassium iodide solution. When ozone
passed through the solution as indicated by the formation
of iodine, the reaction was stopped and the aqueous solution
was extracted with several portions of methylene chloride.
The extracts were concentrated by evaporation of the solvent
and analyzed by g.l.c. and NMR. The aqueous layer was analyzed
for sulfuric, sulfinic and methanesulfonic acids.
Dimethyl disulfide was ozonized by a similar method.
Because the disulfide is not very soluble in water, some of
the reactions were run on suspensions of the disulfide in
27
-------
water. Methylene chloride extracts of the resulting solutions
were analyzed by g.l.c. Evaporation of the methylene chloride
yielded a compound that was distilled under reduced pressure.
The compound was characterized by NMR and IR spectra. The
aqueous portion was analyzed for sulfuric, sulfinic and methane-
sulfonic acids.
Ozone Solutions
During the ozonations using gas streams to introduce
ozone, it was found that large quantities of methanethiol
or dimethyl disulfide were lost in the gas streams. There-
fore, in order to obtain quantitative results solutions of
the thiol or disulfide were added to an ozone solution.
About 4000 g of water was weighed and cooled to 0°C,
for use in the preparation of an ozone solution. Both ozone-
oxygen and ozone-nitrogen streams were employed to make the
solution. Methanethiol solutions prepared as previously
described were made while the ozone solution was being pre-
pared. Because of the instability of the solutions, the
following procedure was carried out as rapidly as possible.
An aliquot of the methanethiol solution was added to an iodine
solution for analysis of the thiol. The initial ozone deter-
mination was started by adding an aliquot of the ozone solution
to a dilute potassium iodide solution. Then an aliquot
of the thiol was added to the ozone solution and stirred
slowly for approximately a minute to insure complete reaction
and mixing. Two aliquots of the reaction mixture were then
added to potassium iodide solutions for determination of the
final ozone concentration. In the initial experiments an
aliquot of the reaction mixture was then used to determine
the methanesulfonic acid. For reasons described later, it was
necessary to remove the residual ozone before titrating. The
ozone was removed by bubbling with nitrogen until a small
sample of the reaction mixture would give no color when
28
-------
added to a potassium Iodide-starch solution. The quantitative
work on dimethyl disulflde was done using the same procedure.
The disulfide was used as a primary standard in making up
solutions and was not analyzed.
29
-------
SECTION V
DISCUSSION AND RESULTS
SULFUR-DIOXIDE OZONE
The various species of sulfur dioxide in aqueous
solution are shown in equation 1.
H?0 . _ ,
(1) S02 J S02 -H20 + H2S03 2 HS03~ + H £ S03~ + H2
It was our initial assumption that each of the species-
sulfite, bisulfite, and sulfurous acid (or aqueous sulfur
dioxide) was capable of reacting with ozone. To distinguish
among these possibilities we measured the rate of the ozonation
reaction at various hydrogen ion concentrations. Kinetics
thus determined yield, for any single pH, an overall rate
constant which may include contributions from one or more
sulfur species. Equations 2-4 below indicate the reactions
which might be occurring.
^(SO )
(2) H2S03 + 03 + 2 H2S04 + 02
H20
kHSO ~
(3) HS03" + 03 -> * HS04
SO =
(4) S03=+ °3 * 3 S04= + °2
H20
Table V shows the mean rate constant (k/tota-i suifur))
30
-------
TABLE V
MEAN RATE CONSTANTS
pH k(total sulfur)
0.
1.
1.
2.
2.
2.
3.
3.
4.
301
620
20
71
11
50
80
12
55
02
2
9
4
1
1
3
4
5
9
1
.24
.82
.31
.02
.63
.03
.57
.97
.47
.74
x
x
X
X
X
X
X
X
X
X
1
1
1
1
1
1
1
1
1
1
o3
0
0
0
0
0
0
0
0
0
3
4
5
5
5
5
5
5
6
+
+
+
+
+
+
+
+
+
.89
.48
.35
.21
.27
.37
.42
.53
.66
16 Determinations
0.59 8.60 x 103 + .23
2.13 1.42 x 105 + .11
3.09 2.13 x 105 + .37
3.74 1.43 x 106 + .28
31
-------
variation with pH at 25.0° and 16.0°. Figure II is a plot
of that data, which shows its regularity in graphical form.
Rate constants were determined using the following
equation:
k = [1/a-b] [ln(b(a-x)/a(b-x))] /t
where k = rate constant
a = initial ozone concentration
b = initial sulfur concentration
x = change in concentration at
t = time
Analysis showed the reaction to be second order, first
order in both ozone and total S(IV) species.
In order to account for the strong dependence of the
rate constant on pH, we have assumed the reaction takes place
at different rates with the three species of S(IV) present due
to the equilibria
H2S03 ? H+ + HS03" K]
HS03" £ H+ + S03"2 K2
The fractional distribution of the species can be calculated
from the pH and the equiIibriurn constants for the above
mentioned equilibria. There seems to be no clear cut choice
for the values to be used with values of K, reported ranging
from 9xlO"3 to 5xlO~2 and K2 ranging from 6xlO"8 to 5xlO~6
at 25°C.
Our calculations reported here are based on the K,values
of Arkhipova et. al.. 35 who report KI values at 35°, 25°,
and 10° as l.OxlO"2, 1.3xlO"2 and l.SxlO"2 respectively.
The K2 values used are those of Teder who gives values of
1.16x10~7 and 9.3xlO"8 for 25° and 60°. These values were
32
-------
20
18
16
14
12
10
8
FIGURE II
DEPENDENCE OF RATE ON pH
33
-------
chosen because they gave a means of calculating values at
temperatures other than those given.
A series of calculations were made using the values K,
= 1.72xl02 and K2 = 1.02xl07. Calculated rate constants
showed about the same variance but absolute values were dif-
ferent due to the different species distribution calculated.
The sulfur species concentrations were converted into
fractions, leading to the equations shown in Table VI. It
was obvious that k^n was considerably smaller than the
other specific rate constants and adds essentially nothing
to the overall rate of reaction at hydrogen ion concentrations
lower than .1 M. Its fraction of the total sulfur is there-
fore not included in equations 4-10.
Equations 3-10 in Table VII were used to determine
k
3 an SO^ (computer program using each equation, one
at a time, with every other equation, and solving for the
rate constants). Those values were then substituted into
equations 1 and 2 to determine the value of kcn . Similar
2
procedures were followed for the data determines at 16°.
Table VII shows the rate constants thus determined.
Arrheneus treatment reads to the following formulas:
d[HSO ~] ,, 11,600
1 "
exp " - [HS03~] [03-]
dt RT
d[SO~2] 10,500 9
— - 1017exp tS03-2] C03]
dt RT
34
-------
TABLE VI
EQUATIONS FOR CALCULATIONS OF SPECIFIC RATE CONSTANTS
General Equations
k = [Fraction HS03]k] + [Fraction S03]k2 + [Fraction S02]k3
25.0° Runs
1) -.301 2.24 x IO3 = .0065 kHSO* + 3.8 x IO"10 k$0= + .993 k$0
3 «3 t
2) 0.620 9.82 x IO3 = .051 kHSQ- + 2.48 x 10"8 k$0= +
8)
9)
0) 4.02 1.74 x 106 = .994 kH$()- + 1.20 x 10
-3
'3 DU3
35
o *
3) 1.20 4.31 x 104 = .171 kHSQ- + 3.14 x IO"7 k$0= + .
3 <5
4) 1.71 1.02 x 105 • .402 kHSQ- + 2.39 x 10'6 ks(J.
3 3
5) 2.11 1.63 x IO5 = .627 kHSQ- + 9.38 x 10"6 k$0=
3 «5
6) 2.50 3.03 x 105 * .805 kHSQ. + 2<95 x 1Q-5 ^^
«5
7) 2.80 4.57 x 105 - .891 kHSQ- + 6.54 x 10"5 k$0=
v »
3.12 5.97 x 105 = .946 kHSQ- + 1.44 x 10"4 k$(J*
3 .
-------
TABLE VI (cont'd)
16.0° Runs
11) 0.59 8.6 x 103 = .058 k, + 3.15 x 10"8 k? + .94 k
£
12) 2.13 1.42 x 105 = .683 k, + 2.76 x 10~5 k2 + .32 k
13) 3.09 2.13 x 105 = .951 k] + 1.64 x 10"4 k£
14) 3.74 1.43 x 106 = .989 k] + 7.61 x 10"4 k
36
-------
TABLE VII
SPECIFIC RATE CONSTANTS
j. O c"
at 25
= 3.1 + 1.1 x 105 M/lsec
k n= = 2.2 + 1.6 x 109 M/lsec
S03
k = 5.9 ± x 102 M/lsec
so2
at 16°
kHSQ- = 1.71 ± .3 x 105 M/lsec
kSO= =1-3 ± .7 x 109 M/lsec
k =0 M/lsec
Kso2
37
-------
from which the reaction rate at any temperature can be
calculated. The overall reaction rate will be the sum of the
two rates determined by calculating the concentration of
each species from the equilibrium constants and the pH.
The contribution of the SO,, • H,,0 species to the rate
is small at all pH of interest in the atmosphere. Our value
_ 2
of less than 10 x kuc ~ would make the contribution of
this species less than 1% of the total rate at all pH values
more than 1.7.
_2
Conversely, the SO, species becomes a greater than
1% contributor when its fraction becomes greater than about
-6 4
10 since its rate constant is about 10 x that of the HSO.,
species. This occurs at about the same pH (1.7) that the SO,,
species becomes unimportant. Thus below pH 1.7 we can ignore
sulfite contribution and above this pH we can ignore the SOp
contribution to the rate. At pH > 4, the reaction is almost
_?
entirely due to S03
Because of the extreme dependence of rate on pH
(rate approximately doubles for 0.5 pH change), the method
of dividing the overall rate constant into the specific rate
constants for the different species is valuable. The distri-
bution among species is a function only of the H concentration
and the ionization constants of the species involved so our
rate constants can be used to determine an overall rate
constant at any desired ph.
There were experimental difficulties in working at the
hydrogen ion concentrations of atmospheric interest. The
problems, were:
(1) Increasing the pH decreases the amount of sulfur
dioxide in solution (bisulfite becoming the
common sulfur species present under our conditions).
This meant that either we increase total sulfur,
in order to have relatively constant absorption
38
-------
of sulfur dioxide for the determination of initial concen-
trations or work at lower accuracies with respect to all
concentrations. We chose to keep relatively high sulfur
dioxide concentrations. In retrospect (see 2&3) our
choice was wrong.
(2) The rates of ozonation increased dramatically with
increasing pH. This meant that at higher pH values our
kinetic results started to be limited by the mixing time of
the stopped flow apparatus. For the 16° runs concentrations
of both ozone and total sulfur were lowered for higher pH
runs, but the determinations above pH 3 at 25° are suspect.
Because of that (after the termination of the grant)
we lowered the concentration of ozone by a factor of 20 and
the concentration of total sulfur by a factor of 80 and did
one more determination at pH 3.22. The observed rate
constant (1.79x10 ) was almost three times larger than what
we would have expected from extrapolation on Figure n.
The result by itself does not indicate whether we had pushed
the limits of the stopped flow apparatus previously or whether
our decreased accuracy of measuring total sulfur was responsible,
In fact we think it likely that a third factor (3 below) may
have been important in both that experiment and in all
kinetic determinations at higher pH.
(3) Our basic assumption has been that equilibrium occurs
in the conversion of one sulfur species to another. Using
the rate constants from equation 5 it can be shown that equil-
ibrium would be established rapidly enough so that ozonation
can occur with each species at most concentrations used.
However, at high pH and with very high concentrations of
total sulfur and ozone the rate of formation of sulfite from
sulfurous acid is only about ten times that of its dis-
appearance from reaction with ozone. This seems not to
39
-------
be a particularly bad problem since at such acidities most
of the total sulfur is already in the bisulfite form. More
serious is our original assumption that the rate constants
for the second step of the ionization
k2 = 104
HSO " t H+ + SO"
11 ^
k_2 = 10''
has the values shown. We have found no mention of these
rate constants in the literature and have assumed them from
a) knowing the equilibrium constant and b) guessing that
the recombination of a proton with sulfite ion might have
about the same rate constant as the similar reaction between
sulfate ion and a proton (a diffusion controlled reaction
rate). At pH above about 2, the SO., rate is important.
2
The above assumed rates would indicate that SO, is being
used up more rapidly than it can be replenished by HS03~
dissociation. This would result in a trend similar- to that
observed of an apparent relative decrease in the constant
with the increase in pH.
Environmental Implications
The reason for determining the kinetics of the reaction
between ozone and sulfur dioxide in solution was to discover
whether such a reaction might take place under normal atmo-
sphere conditions. While this study was in progress, Penkett
determined the rate of the reaction between bisulfite ion and
ozone and reported the following results.
a) The reaction is rapid, having a second order rate
of
dO. _
—- [3.32+ .11 x 10V sec "'] [Oj [HSO ~]
dt " J
at 9.6° and at a pH of 4.65.
40
-------
b) Using reasonable atmospheric conditions [SO,, in
atmosphere = 0.007 ppm, 0- - 0.05 ppm, clouds
contain 0.1 to 1 g liquid water per cubic meter,
bisulfite concentration in solution = 5 x 10" ]
he calculates an oxidation rate of 12.6% hr .
He notes that his is 70 times the rate of disappearance
of sulfur dioxide via oxygen.
Our results are in conflict with those of Penkett on
several counts. First we note that the pH dependence of
the reaction is such that there seems to be little doubt
that Penkett measured the rate of the sulfite ion-ozone
reaction as a major part of his rate constant. Since his
ozone solution was unbuffered and at a pH of 4 and his bi-
- 4 - 5
sulfite concentration varied between 10 and 10 the exact
pH (and therefore the exact sulfite ion concentration) must
have been slightly variable. We have retabulated his rate
constant variations with bisulfite ion concentration in Table
IX.
Table VIII
Rate Constants Ozone-bisulfite (Penkett)
cone HS00 (M) k
^
10"5 5.7 x 105
2.5 x 10"5 3.48 x 105
5.0 x 10"5 3.46 x 105
10"4 3.27 x 105
If Penkett's initial pH were 4.65, as he suggests,
then the normal prediction would be that the solution w.th
least buffering capacity (the 10" solution) would be at
the lowest pH when mixed with the ozone solution at pH 4.
41
-------
Our initial prediction would have been that rates would
increase with increasing bisulfite ion concentration since
the highest pH solution would have the highest sulfite
ion concentration. The trend of Penkett's data is clearly
in the other direction. The highest rate is found with the
lowest initial concentration of bisulfite.
However when the major discrepancy of our work with
Penkett's is considered, a possible explanation for his
upward trend in rates with decreasing bisulfite may be
given. Specifically, our rate equation predicts that the
second order rate constant for Penkett's conditions should
have been 5.5 x 106, using HSQ,"= 10"5 M. His value of 5.7
5
x 10 would therefore be too low by a factor of about 10,
while his "average" rate constant (3.32 x 10 ) is about 15
times lower than what we would predict.
We believe that Penkett's upward trend in rates with
decreasing bisulfite concentrations, our own similar finding,
and the fact that our predicted rate constants are consider-
ably higher than his experimentally determined values, have
the same cause. Specifically, sulfite can react with ozone
only if it is formed rapidly enough under equilibrium con-
ditions to do so.
For purposes of comparison, let us assume the same
concentrations as Penkett did: S09 at 0.007 ppm, ozone at
3
0.05 ppm, .1 gm per m liquid droplets of water, pH = 5.0,
and a temperature of 10°C. Using the activation energies
obtained in this work we can extrapolate values for k,,cg-
and kQn= of 1 .
^3 i
respectively.
Using the
1 x 105 (M/l
values of KH
r1 sec'1
Mso2 (
Pso2 (
and
aq)
g)
7.4 x 108 (M/l)"
= 2.20 and
42
-------
(H+) (HSO")
K, = = .0184 from the work of Johnstone and Leppla
(SO- (aq) )
£ c
(1934) we arrive at a HSO" concentration of 4 x 10 M
= -7
and a SO, concentration of about 4 x 10 M. and an ozone
-9
cone, of 10 M. Therefore,
- dS(IV) = (1.1 x 105 x 4 x 10"5 + 7.4 x 108 x 4 x 10"7)xlO"9
dt 7
= 3. x 10"' (M/l) sec -1
3 -4
For each M we have 1 x 10 1 of solution so we now obtain
3 x 10" moles/m /sec reacted or 1 x 10" moles/m hr or
6.4 x 10"6 g/m3 hr.
The starting concentration of .007 ppm gives 20 mg m~
so 6.4 = .32 or 32% per hr is reacted. Penkett arrived at
20
a figure of 12.6% per hr. This difference is partially due to
the apparently different values used for K,, and K, (He
calculated a concentration of HSO^ = 1 x 10 ).
Another possible contributing factor may lie in the precision
of our activation energy and the consequent uncertainty of
the extrapolated values for the specific rate constant. All
of the above calculations assume that solution is rapid enough
that equilibrium between the various phases and species is main-
tained.
Comparing the above rate with rates published for other
oxidation modes, it must be concluded that, for the cited
conditions (liquid droplets, presence of ozone at a constant
concentration) ozone caused oxidation is an important contri-
bution to the oxidation of S(IV) in the atmosphere. For exam-
•i /•
pie, Brimblecombe and Spedding using approximately the same
total S(IV) and H20 (1) values that we assume, calculate the
43
-------
removal of about 3% per day by the Fe (III) catalysed
37
oxidation by atmospheric oxygen, Cheng et. aj_. calculate
a 2%/hr decrease in sulfur dioxide in natural fog in the presence
38
of manganese salts and Sidebottom et. al. show a 1.9%/hr
loss of sulfur dioxide photochemically at high humidity.
44
-------
METHANETHIOL-OZONE
Methanethiol-General Results
When an ozone oxygen stream was passed into a 0°C
aqueous solution of methanethiol, a rapid reaction occurred.
Quantitative analysis was not possible using this system
because large amounts of methanethiol were carried away in the
oxygen stream. Cold traps were not effective in trapping
all of the escaping thiol, but silver ions were found to
precipitate all of it. Since the determination of the
amount of thiol precipitated was not practical, an indirect
method of analysis was attempted. A known amount of silver
nitrate was used to trap the escaping thiol and the excess
silver was titrated with potassium thiocynate. Unfortunately,
absorption of silver ions on the precipitated mercaptide
obscured the end point. Therefore, the reaction was usually
stopped when the thiol concentration was so low that some
of the ozone passed through the solution without reacting.
A strongly acidic solution resulted from the reaction;
however, qualitative tests for sulfuric and methanesulfinic
acids were negative. The acidic product was isolated and
it was identified as methanesulfonic acid by conversion to
methane-sulfonyl chloride. Extraction with methylene chloride
revealed that other products were also present. Dimethyl
disulfide, methyl methanethiosulfonate and methyl methane-
thiosulfinate were identified as minor constituents. The
disulfide and thiosulfonate were identified by coinjection
gas chromatography and their mass spectra. The thiosulfinate
was shown to be present by gas chromatography and nuclear
magnetic resonance resonance spectra. Continued ozonation
resulted in the formation of sulfuric acid. The equations
45
-------
below summarize the chemical processes:
CH3SH
slow
CH3SO 'SCH3
CH3S02SCH3 Xtrace)
The stoichiometry of this reaction was studies by
adding known thiol solutions to known ozone solutions.
The results obtained are given in Table X. The results
indicate a stoichiometry of slightly less than 2 moles of
ozone per mole of thiol. No difference in stoichiometry
was found, whether ozone-oxygen or ozone-nitrogen
streams were used to prepare the ozone solutions.
Titration of the resulting sulfonic acid was not straight-
forward. When a sample was titrated immediately following
reaction, the titration curve contained two inflection points.
The first inflection corresponded closely to quantitative
conversion of the thiol to sulfonic acid. If the solution
was allowed to stand several days before it was titrated,
a single inflection point having a value equal to that of
the second inflection in the original solution was obtained.
It was found that if the excess ozone was removed before titra-
ting, only one inflection point was obtained and it correspond-
ed to near quantitative conversion of the thiol to sulfonic
acid. The yield of sulfonic acid was taken at the first inflec-
tion point was obtained and it corresponded to near quantitative
conversion of the thiol to sulfonic acid. The yield of sul-
fonic acid was taken at the first inflection point or from
a titration in which the excess ozone had been removed. The
yields obtained are also given in Table X.
The rate of the reaction appeared to be very fast.
46
-------
TABLE IX
Run Moles of 03 Moles of 03/ Moles of %
Consumed CH3SH Added CH-jSH CH3$03H Produced Yield
1
2
3
4
5
7
/
8
1.24 x 10"3
1.18 x 10'3
#9.7 x 10"4
#4.3 x 10"4
1.03 x 10'3
Q £ y y 10
J • U A A 1 U
#1.82 x 10'3
6.
6.
5.
2.
5.
O •
1
1 •
1.
30 x
84 x
58 x
38 x
97 x
7 1 v
/ O A
fiK y
U J A
03 x
TO'4
TO'4
ID'4
TO'4
TO'4
IO-4
1 U
TO'4
1 U
TO'3
1
1
1
1
1
1
1
.
1
.97
.75
.74
.80
.73
fifi
• o o
— _
.76
6.
6.
5.
2.
7.
1.
1.
36 x
87 x
54 x
74 x
40 x
71 x
02 x
10
10
10
10
10
10
10
-4
-4
-4
-4
-4
-4
-3
101
100
99.3
115
124
*103
98.5
1.09 x 10"3 *106
# - Solution made with ozone-nitrogen stream.
* - Yield determined after removal of excess ozone
47
-------
Observation of the decrease in ozone was followed spectrophoto-
metrically at 285 nm using the stopped flow apparatus. At
initial thiol concentrations of 7 x 10 M. and ozone concen-
trations of 3 x 10 M the reaction was over in about 0.1
seconds.
The experimental titration curves show a double hump
pattern obtained when an aliquot of the reaction mixture
is titrated in the presence of ozone. However if the ozone
is removed a standard titration curve is obtainable.
This pattern, although puzzling at first, proved to have a
simple explanation. In basic solutions, ozone will react
with methane sulfonic acid to produce sulfuric acid much
faster than in acidic solutions. In acidic solutions, the
reaction proceeds very slowly. That is why letting the solu-
tion stand for several days results in a single hump pattern
corresponding to the second hump in a titration with ozone
present. As the equivalence point is reached during titra-
tions with ozone present, the pH rises allowing sulfuric
acid to be produced at a reasonable rate. The pH then remains
relatively constant until the ozone has all reacted. Tests
with methanesulfonic acid showed the same behavior as the
reaction mixtures. Other tests demonstrated that sulfate
was formed when ozone was added to basic solutions of methane-
sul foni c acid.
Methanethiol-Errors
Any quantitative measurement is subject to some error.
The results given in Table X exhibit some random error.
The stoichiometry from run #1 and the yield from run #5
are more than two standard deviations from the mean and,
therefore, have been dropped from the respective calculations.
48
-------
Doing this one obtains a stoichiometry of 1.74 + 0.04
and a yield of 103% + 6%. The limits of error given are
one standard deviation.
This treatment of random errors does not include any
systematic errors. Although efforts were made to reduce any
error of this type, several sources of possible error exist.
The solutions used in this study were all unstable with
respect to loss of the solute into the air. In order to
minimize this type of error, initial concentrations were
measured, the reaction was carried out and final concentra-
taions were measured as rapidly as possible. Ozone is also
known to undergo thermal decomposition, but this is a slow
reaction compared to the time between samplings (about one
minute).
Because of the limited solubility of ozone in water,
the solutions used were very dilute. Final ozone concentra-
- 5 -4
tions ranged from 5 x 10 to 3 x 10 M. At these concentra-
tions, end points were not extremely sharp. This was par-
ticularly a problem in the sulfonic acid determinations where
the steep portion of the titration curve generally occurred
in 1/4 to 1/2 mill il Her.
Methanethiol- Scientific Implications
The reaction of ozone with methanethiol has been shown
to produce methanesulfonic acid and eventually sulfuric acid.
The conversion of the sulfonic acid to sulfuric acid is
very slow and, for practical purposes, the reaction stops
at the sulfonic acid. When insufficient ozone was added
directly to a solution of thiol, dimethyl disulfide, methyl
methanethiolsulfinate and methyl methanethiolsulfonate
were also detected. The quantitative results seem to indicate
that these products were not formed in the reactions carried out
49
-------
by mixing solutions. This is reasonable because, in the
presence of excess ozone, any disulfide and thiolsulfinate
would have been oxidized further. However, Barnard has
reported that thiolsulfonates are resistant to ozonation.
No minor products were detected from these reactions; but
at such low concentrations, they may have escaped detection.
These results are not surprising and are similar to
the results reported for other thiols. A thiolsulfinate, how-
ever, has not previously been reported as a minor product
or intermediate in the ozonation of a thiol. The reason
for this probably involved a difference in the mechanism
of disulfide ozonations in aqueous and organic solvents.
Previous workers using organic solvents have not
found thiolsul finates to be products of long-lived intermediates
in the ozone-disulfide reaction. This evidence, although
it seems to indicate that the thiolsulfinate is formed from
the disulfide, is not sufficient for proof.
Because thiols are readily converted to disulfides, it
might be expected that the ozonation of thiols proceeds via
disulfides. This does not appear to be the case for methanethiol
If the reaction involved the disulfide, the overall rate would
be controlled by the rate of ozonation of the disulfide or
the rate of conversion of the thiol to the disulfide, whichever
is slower. But methanethiol was observed to react faster
than dimethyl disulfide with ozone. Therefore, it must
be concluded that disulfide production is only a minor side
reaction.
The stoichiometry found is similar to those obtained
by Barnard17 (1.6 03:SH) and Erickson (1.6 03:iso-BuSH)
when the ozone-oxygen streams were passed into the thiol
solutions.
50
-------
Methanethiol-Envi ronmental Imp!ications
The reaction product (CHoSCKH) from the ozone-methanethiol
reaction is odorless, nonvolatile, and water soluble. These
are all positive attributes that could be taken advantage
of in designing pollution control equipment for Kraft pulp
mills. Although the rate data obtained are in no way complete,
it appears that the reaction is fast enough so that no major
technological difficulties should be encountered in equipment
design. These are some difficulties in using aqueous phase
ozonations for pollution control. However, use of ozonized
water in the water scrubbers for stack gases or direct gas
phase ozonation might be possible.
The importance of this reaction in the atmosphere is
hard to judge. In regions where ozone and thiols are present
in the atmosphere together, some reaction no doubt takes place,
However, these reactions are probably occurring in the gas
phase for the most part. In plumes from pulp mills, where
suspended water droplets are in high concentration, some
aqueous phase reaction may occur if ozone is present.
The situation in the upper atmosphere is similar; if
thiols and ozone are both present, some reaction probably
occurs, but in the gas phase. Two things should be kept
in mind when dealing with the upper atmosphere: 1) it has not
been shown that thiols are present there, and 2) other path-
ways involving the high energy species would compete for any
thiol present.
DIMETHYL DISULFIDE-OZONE
Dimethyl Disul fide-General Results
When an aqueous solution of dimethyl disulfide was
51
-------
ozonized by passing an ozone-oxygen stream into the solution,
the major product formed was methanesulfonic acid. The acid
was identified by converting it to methanesulfonyl chloride
followed by coinjection gas chromatography and determination
of its mass spectrum. Methyl methanethiolsulfinate was observed
to form in high yield during the beginning of the reaction.
In one experiment 0.0056 moles of ozone was added to 0.011
moles of dimethyl disulfide dissolved in two liters of water.
The only product detected by NMR spectra of a methylene
chloride extract was methyl methanethiolsulfinate. The
final pH of the reaction mixture was 5.3. In some runs a small
amount of methyl methanethiolsulfonate was also formed.
No sulfinic acids were detected. This is summarized in equation
°3
CH3SSCH3 + 03 -> CH3SO -SCH3 -> CH3$03H
The stoichiometry and yields found for the disulfide reaction
are given in Table XI. All titrations of the sulfonic acids
were performed after the removal of excess ozone from the
solution.
The rate of this reaction appeared to be slower than
-4
that of the thiol. At initial concentrations of 1 x 10
M CH3SSCH3 and 6 x 10" M 03 the reaction was observed to be
complete in about one half minute.
Dimethyl Pisulfide-Errors
The mean stoichiometry and yield found were 3.95 +
0.45 and 94.8% + 7.1%. Eliminating runs more than two standard
deviations away from the mean gives 3.89 + 0.07 and 95.6%
+ 4.0%.
This reaction was carried out using the same procedure
as the thiol reaction, and the limitations of the method
are the same. One additional problem was encountered,
52
-------
TABLE X
Run Moles of 0., Moles of 0, / Moles of %
3 0
Consumed CH3SSCH3 CH3SSCH3 CH3S03H Yield
Added Produced
1
2
3
4
c;
6
*2.
*2.
2.
2.
2.
1.
24 x
17 x
68 x
15 x
21 x
90 x
10"
10"
10"
10"
10"
10"
q
o
•3
o
3
3
•3
*}
q
o
5.63
5.63
5.63
5.63
5.63
5.63
x 10
x 10
x 10
x 10
x 10
x 10
-4
*T
-4
~ *T
-4
-4
-4
*T
-A
*T
3
3
4
3
3
3
.97
.84
.77
.82
.92
.38
1
1
1
1
9
1
.16 x
.13 x
.10 x
.04 x
.37 x
.04 x
_ q
10 J
_ q
10 6
10"3
10"3
-4
10 4
q
10 J
103
100
98.1
92.2 -
83.2
92.2
- solution made with an ozone-oxygen stream.
53
-------
however. The time necessary for completion of the reaction
was not considerably less than the time allowed for reacting.
Although enough time for complete reaction was generally
allowed, it is remotely possible that some reactions were
stopped before completion. Longer reactions times were
avoided because of the instability of the solution.
Dimethyl Disulfide-Scientific Implications
There appears to be a difference in yield of sulfonic
acid between reactions run with solutions made from ozone-
oxygen and ozone-nitrogen streams. Considering the limited
amount of data and the large degree of scatter, it is possible
that this apparent difference is not real. It is probably
best to conclude that ozone will convert dimethyl disulfide
to methanesulfonic acid in high yield.
Sulfonic acids have not previously been reported as
the products from the reactions of disulfides with ozone,
but the formation of sulfonic anhydrides was observed in
nonaqueous solvents by Barnard . Although water must be
involved at some point in the reaction sequence, the reaction
does not necessarily follow the mechanism proposed by Barnard
with the addition of a hydorysis step at the end. In fact,
the results of the investigation indicate that a different
mechanism is in operation. In trying to elucidate the mechanism
of the reaction, Barnard conducted experiments in which
insufficient ozone for complete conversion was used. Only
the normal products of complete ozonation were found.
In this study, however, high yields of methyl methanethiol-
sulfinate were isolated when insufficient quantities of ozone
were used. Apparently in water, the first step of the conversion
to sulfonic acid is the oxidation of the disulfide to the
thiolsulfinate. Although no sulfonic acid was formed early
in the reaction, quantitative data is needed to show that
54
-------
another reaction pathway going through an undetected interme-
diate was not in operation. The small amounts of thiol-
sulfonate that were detected in some runs probably arose through
minor side reactions. Known reactions that could possible
be the source are:
1) the disproportionation of thiolsulfinate to disulfide
and thiolsulfonate, and
2) the autoxidation of thiolsulfinate. It is also
possible that the thiolsulfonate is a direct
ozonation product.
Dimethyl Disulfide-Environmental Implications
Because methanesulfonic acid is the major product of
this reaction as well as from methanethiol, the same advantages
(odorless, nonvolatile, and water soluble) in designing
pollution control equipment are gained. The slower reaction
rate would make removal of the disulfide more difficult.
The opportunity for reactions in the atmosphere are the
same as those for the thiol.
DIMETHYL SULFIDE-OZONE
Over a year's work went into the building of gas phase
kinetic system, and many months of obtaining rate data follow-
ed. Although the data is available , it is meaningless.
Our original system apparently measured the rate of reaction
of ozone with dimethyl sulfide which had coo ed out in the
freeze out trap. Near the end of the research period several
new reaction chambers were built and it appeared as if the
reaction was suitable for study.
The last of the reaction chambers was built from non-
precision bore 20 mm (OD) tubing and was 120 cm long (Volume =
305 ml). The bend into the KI sampler was considerably
less restricted than earlier models. The following data was
determined.
55
-------
Run
Flow (1/min)
Order uncorrected = 1.314
corr. coef. = .995
Run
Flow (1/rnin)
M(03)xlO(
M(DMS)xlO'
d(0-)/dtxlO'
1
2
3
4
5
6
7
8
1
1
1
1
1
1
1
1
.260
.278
.250
.261
.273
.252
.259
.275
9.
11.
7.
2.
2.
6.
7.
8.
049
362
054
481
724
702
425
797
1 .646
1 .912
1.103
.2779
.3241
1.242
1.273
1 .498
d(03)/dtx!0-
1
2
3
4
5
6
7
1.265
1.261
1.274
1 .267
1 .272
1 .275
1 .260
1.
•
•
3.
.
2.
2.
462
9874
9773
499
6657
276
303
1.1967
.7842
.7080
1.7399
.5467
1.498
1.646
Order uncorrected = .757
corr. coef. = .974
Correctee Ozone Order = 1.398
Corr. Coef. = .990
Corrected DMS Order = .861
Corr. Coef. = .969
* This order was computed assuming a 1:1 stoichiometry
for ozone and DMS. If the ozone order is assumed to be 1.5
and the DMS to be 1.0, the rate constant at 25° is approximately
4 10
moles/1 min. This might be compared to the rate
56
-------
22
found by Hales for the gas phase reaction between hydrogen
sulfide and ozone of
Repetition of the method at 35° yielded inconsistent
results. Specifically ozone concentrations became highly
erratic. Whether these problems arose from a faulty ozonator
or the method itself is still unknown.
57
-------
REFERENCES
1. Urone, P. and Schroeder, W.H., "S02 in the Atmosphere:
A Wealth of Monitoring Data, but Few Reaction Rate Studies"
Env. Sci. and Tech. 3, 436-445, 1969.
2. Bufalini, M., "Oxidation of Sulfur Dioxide in Polluted
Atmospheres - A Review" Env. Sci. and Tech. 5, 685-700,
1971.
3. Kellog, W.W., Cadle, R.D., Allen, E.R., Lazrus, A.L. and
Martell, E.A., "The Sulfur Cycle" Science 175, 587-596,
1972.
4. Williams M.D. "Review of Selected Literature on Sulfur
Dioxide to Sulfate Aerosol Conversion".
5. Gregor, I.K. and Martin, R.L., Reaction Between Ozonized
Oxygen and H9S in the Gaseous Phase Austral J. Chem.,
14, 462, 196T.
6. Cadle, R.D. and Ledford, M., The Reaction of Ozone with
Hydrogen Sulfide Air and Water Pollut. Int. J. 10,25,
1966.
7. Pitts, J.N., Environmental Appraisal: Oxidants, Hydro-
carbons, and Oxides of Nitrogen, J. Air Pollut. Control
Assoc., 19, 658, 1969.
8. Cadle, R.D.Discuss ion, ibid. 19, 668, 1969.
9. Espenson, J.H. and Taube, H., Tracer Experiments with
Ozone as Oxidizing Agent in aqueous solution; Inorg.
Chem., 4, 704, 1965.
10. Cox, R.A. and Penkett, S.A., Oxidation of Atmospheric
Sulfur Dioxide by Products of the ozone-olefin reaction
Nature, 230, 320, 1971.
11. Penkett, S.A., "Oxidation of S02 and other Atmospheric
Gases by Ozone in Aqueous Solution" Nature, 240, 105-6,
1972.
12. Akamatsu, I., Kamishima, H. and Kimura, Y., Kami-pa
Gikyoshi, 1968, 22, 406-410, Deodorization of Exhaust
Gas in Kraft Pulping I Formation of Malodorous Components
in Kraft pulping.
13. Maneck, M., Das Braunkohlen-Archiv.40, 61 (1933). Desul-
furization of brown-coal tar distillates.
58
-------
14. Bohme, H. and Fischer, H., Ber. . 75B.J310 (1942).
Uber die Einwirkung von Ozon auf ThToather.
15. Boer, H. and Kooyman, E.G., Analy. C h i m. Acta, 5, 550
(1951). The Use of ozone as a ti trimetric agent for
the determination of definic unsaturation.
16. Bateman, L., and Cunneen, J.I., J. Chem. Soc, 1596
(1955). Oxidation of Organic Sulfides III-
17. Barnard, D., J. Chem. Soc. . 4547 (1957). Oxidation of
Organic Sulfides IX The Reaction of Ozone with Organic
Sulfur Compounds-
18. Maggiolo, A. and Blair, E.A., "Ozone Chemistry and
Technology" Advances in Chemistry Series No. 21,
American Chemical Society, pp. 200-01.
19. Homer, L., Schaefer, H. and Ludwig, W., Ber., 91, 75
(1958). Ozonisierung von tertiaren Aminen (Phosphinen
Und Arsinen) sowie Thioathern und Disulfiden.
20. Kimura, H. Huges L.J., McMinn Jr., T.D. and Burleson,
J.C., C.A.. 60, 6751 (1964).U.S. 3, 114, 775 (Cl. 260-607)
21. Akamatsu I, Kami-pa Gikyoshi 22 (4), 200 (1968) and
Akamatsu I, Kamishima, Kami-pa Gikyoshi 20 (w)
556 (1966).
22. Hales, J., Doctorate Thesis, University of Michigan,pp. 69-
70, 1968.
23. Asinger, F., Ebeneder, F. and Richter, G., J. Prakt. Chem.,
2,203 (1955). Products of the joint action of sulfur
dioxide and chlorine on alphatic hydrocarbons in U.V.
light VI.
24. Erickson, R.E., unpublished research.
25. Kirchner, K., Kastenhuber, H. and Biering, L., Chem. Ing.
Tech. , 43, 626 (1971). Kinetics of the reaction between
mercaptan and ozone in the ppm range.
26. Karchmer, J.H., "The Analytical Chemistry of Sulfur
and its Compounds", Part II, Wiley Interscience, New
York, 1972, p. 764.
59
-------
27. Weast, R.C., et. al . , "Handbook of Chemistry and Physics,"
45th Ed., Chemical Rubber Co., Cleveland, Ohio, 1964, p.
D-103.
28. Jose, J., Philippe, R. and Clechet, P., Bui 1 , Soc. Chim.
France 2860 (1971). Etude des equilibres liquide-
vapeur par mesure de la pression totale -DMSO.
29. Boyd, A.M., Willis, C. and Cyr, R., Anal. Chem.. 42.
670 (1970). New Determination of Stoichiometry of the
lodometric Method For Ozone Analysis at pH 7.0.
30. Hodgeson, J.A., Baumgardner, R.E., Martin, B.E. and
Rechme, K.A., Anl. Chem 43, 1123 (1971).
31. Kopczynski, S.L. and Bufalini, J.J., Anl . Chem., 43,
1126 (1971). Some Observations on Stoichiometry of
lodometric Analyses of Ozone at pH 7.0.
32. Ref. 25, p.481.
33. Feigl, F., "Spot Tests in Inorganic Analysis,"
Elsevier Publishing Company, New York, 1958, pp. 314-
315.
34. Ref. 25, Part II, pp. 767 and 803.
35. Arkhipova, G.P., Mischenko, K.P., and Flies, I.E.
"Equilibrium during disolution of gaseous sulfur dioxide."
36. Brimblecombe P, and Spedding D.J. Atmospheric Environment
8, 937 (1974) The Catalytic Oxidation of Mfcromolar
Aqueous Sulfur Dioxide.
37. Cheng R., Com M., and Frohliger J., Atmospheric Environ-
ment 5,987-1008, 1971, Contribution to the Reaction Kinetics
of Water Soluble Aerosols and S02 in Air at PPM Concentrations
38. Sidebottom H.W., Badcock C.C., Jackson G.E., Calvert J.C.,
Remhardt G.W., and Damon E.K., Envi ron. Sci. Technol.6, 72-79,
1972, Photoxidation of Sulfur Dioxide.
39. Moody R., Thesis, U of Montana, 1973, "Kinetics of the Ozon-
ation of Dimethyl Sulfide in the Gas Phase" 63 pp.
60
-------
TECHNICAL REPORT DATA
(Please read Instructions on the reverse before completing)
1 Ttff-W/3-76-089
2.
3. RECIPIENT'S ACCESSIOWNO.
4. TITLE ANDSUBTITLE
REACTION KINETICS OF OZONE WITH SULFUR COMPOUNDS
5. REPORT DATE
August 1976
6. PERFORMING ORGANIZATION CODE
7 AUTHOR(S)
Ronald E. Erlckson and Leland M. Yates
8. PERFORMING ORGANIZATION REPORT NO.
9. PERFORMING ORGANIZATION NAME AND ADDRESS
Chemistry Department
University of Montana
Missoula, Montana 59801
10. PROGRAM ELEMENT NO.
1AA008
11. CONTRACT/GRANT NO.
R-800655
12. SPONSORING AGENCY NAME AND ADDRESS
13. TYPE OF REPORT AND PERIOD COVERED
Environmental Sciences Research Laboratory
Office of Research and Development
U. S. Environmental Protection Agency
Research Triangle Park, NC 27711
Final Report
14. SPONSORING AGENCY CODE
EPA-ORD
15. SUPPLEMENTARY NOTES
16. ABSTRACT
The reaction between ozone, sulfur dioxide, dimethyl sulfide, methanethiol
and dimethyl disulfide are reported.
The rate of reaction between ozone and dissolved sulfur dioxide is
strongly pH dependent. These results suggest that under some conditions
atmospheric oxidation of sulfur dioxide may involve ozone. Dimethyl
sulfide was found to react rapidly with ozone in the gas phase but
reproducible kinetic data were not obtained. Stoichiometric and yield data
from the reaction between ozone and methanethiol or dimethyl disulfide in
aqueous solution indicate those reactions to be complex mechanistically,
although one product, methane sulfonic acid is predominant.
17.
KEY WORDS AND DOCUMENT ANALYSIS
DESCRIPTORS
b.lDENTIFIERS/OPEN ENDED TERMS
c. cos AT I Field/Group
*Air pollution
*Reaction kinetics
* Ozone
*Sulfur dioxide
*Sulfur organic compounds
13B
07D
07 B
07C
18. DISTRIBUTION STATEMENT
RELEASE TO PUBLIC
19. SECURITY CLASS (ThisReport)
UNCLASSIFIED
21. NO. OF PAGES
69
20. SECURITY CLASS (This page)
UNCLASSIFIED
22. PRICE
EPA Form 2220-1 (9-73)
61
------- |