PB86-216751
Oxygen-18 Study of S02 Oxidation in
Rainwater by Peroxides
Argonne National Lab., IL
Prepared for
Environmental Protection Agency
Research Triangle Park, NC
Jul 86
U.S. Department of
Technical I
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TECHNICAL REPORT DATA
(Please read lauruciions on the reverse before completing)
1. REPORT NO.
EPA/600/3-86/035
3. RECIPIENT'S ACCESSI Of* NO.
2 1 6 y 5 1
4. TI-..E AND SUBTITLE
OXYGEN-13 STUDY OF S02 OXIDATION IN
RAINWATER BY PEROXIDES
b. REPORT DATE
July 1936
6. PERFORMING ORGANIZATION CODE
7. AUTHORIS)
B. D. Holt and R. Kumar
8. PERFORMING ORGANIZATION REPORT NO.
9. PERFORMING ORGANIZATION NAME AND ADDRESS
Chemical Technology Division
Argonne National Laboratory
Argonne, Illinois 60439
10. PROGRAM ELEMENT NO.
CCVN1A/Q2-3058 (FY-86)
11. CONTRACT/GRANT NO.
IAG DW 89930060
. 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 REPORT AND PERIOD COVERED
Final (10/82 - 5/8.6)
14. SPONSOFIIN3 AGENCY CODE
EPA/600/09
15. SUPPLEMENTARY NOTES
16 ABSTRACT
A new analytical method was developed for the determination of oxygen.,
isotope ratios in peroxides in rainwater. In the method, rainwater samples
were quantitatively degassed of dissolved air by a combined treatment of
evacuation, ultrasonic agitation, and nelium sparging (VUS), followed
by a permanganate oxidation of the dissolved peroxide to 02. The 02 was
then quantitatively removed fron the rainwater by the VUS treatment and
converted to C02 for mass spectrometric analysis. Using this method,
14 rainwater samples collected at four sites (Argonne, IL, Research
Triangle Park, NC, Uhiteface Mountain, NY and Dearborn, MI) were analyzed
to determinne the fraction of the sulfate in the samples that was produced
as a result of the aqueous-phase reaction of dissolved S02 with hydrogen
peroxide. It was concluded that 40% or more of the sulfate in the samples
was formed by peroxide oxidation.
17.
KEY WORDS AND DOCUMENT ANALYSIS
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21. NO. OF PAGES
43
RELEASE TO PUBLIC
30. SECURITY CLASS (This page)
illsiri ASSTFTFil
22. PRIC
t
EPA Form 222C-1 (1-7))
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DISCLAIMER
The information in this document has been funded by the
United States Environmental Protection Agency under In-
teragency Agreement DW 030060-01-0 to Argonne National
Laboratory. It has been subject to the Agency's peer and
administrative review, and it has been approved for pub-
lication as an EPA document. Mention of trade names or
commercial products does not constitute endorsement or
recommendation for use.
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ABSTRACT
Oxygen isotope ratio measurements were used to examine the importance of peroxide
oxidation of SO3 to sulfate in the atmosphere. A new analytical method was developed for
the determination of oxygen isotope ratios in peroxides (concentrations in the ppb range)
in rainwater. According to the method, 20-L samples were quantitatively degassed of
dissolved air by a combined treatment of evacuation, ultrasonic agitation, and helium
sparging (VUS), followed by permanganate oxidation of the dissolved peroxide to O2;
the C>2 was then quantitatively removed from the rainwater by the \TJS treatment and
converted to COj for mass spectrometric analysis. Stock solutions of HjOa of various
18O enrichments were prepared by a high-voltage discharge method; and, by using these
solutions to oxidize SOa to SO^~, the following .sotopic relationship was established:
_ = 0.57 «51802- + 0.435180H0 + 8.4%.
This relationship was used to calculate (he <518Ocr.2- formed by peroxide oxidation and
^U4
a similar previously established relationship was used to calculate the 618O<-,.~ 2— formed
by metal-catalyzed aqueous oxidation. These sets of calculated values were compared to
the measured values for evaluation of the estimated fraction of sulfate in rainwater that
was formed by peroxide oxidation, assuming that metal (or carbon) catalysis was the only
other major sulfate formation mechanism. It wns concluded that ~40C7> or more of the
sulfate in summer rains in the northeastern U. S. was formed by peroxide oxidation.
ill
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CONTENTS
ABSTRACT iii
TABLES vi
FIGURES vii
ACKNOWLEDGMENT viii
1. INTRODUCTION 1
2. EXPERIMENTAL . 3
Development of Analytical Method ... 3
Conversion of O2 to CO2 • . .... 7
Removal of Dissolved O2 from Water 9
Oxidation of H2O2 to O2 in Water . . 13
Synthesis of 18O-enriched H2O2 13
Collection of Rain Samples 14
3. RESULTS 16
Recovery of Added H2O2 . 16
Test for Isotopic Interference 16
Effects of Autodecomposition of H2O2 16
Effects of KMnO4 Decomposition 18
Procedure Blank ... .18
Isotopic Relationships 18
<518O and Concentrations of H2O2 in Rainwater 22
4. DISCUSSION 26
6-8O of H2O2 in Rainwater 26
618O of Atmospheric Sulfates: Measured vs. Calculated 27
Fraction of Atmospheric Sulfate Formed by Peroxide Oxidation .... 29
5. SUMMARY 32
REFERENCES 33
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TABLES
Page
1. Concentration and isotopic results for rainwater collected at Argonne
National Laboratory (ANL), Dearborn, MI (DEA), Whiteface Moun-
tain, NY (WFM), and Research Triangle Park, NC (RTP), May-
September, 1985 23
VI
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FIGURES
Number Page
1. Four-step method for conversion of oxygen in H2O2 (ppb range in
water) to CO2 for mass spectrometric analysis ............ 4
2. Analytical train ........................... 5
3. Reaction chamber ....................... 6
4. Recovery of O2 as CO2 after reaction with charcoal in contact with
platinum at various temperatures ................. 8
5. Recovery of CO2 after exposure to charcoal and platinum at various
temperatures ......................... 10
6. Removal of O2 from water by various degassing techniques ........ 12
7. Isotopic effects of water solvent on the <518O of CO2 originating from
H202 ............................. 17
8. Dependence of £l8OjjoQ on 518Oj| Q in a high-voltage discharge
process of formation ........................ 19
9. Dependence of 6 18OC,^.2_ on <518O|j Q in oxidation of SO2 to sulfate . . . 21
10. Measured isotopic data for II2O2, SO^~, and H2O in the samples of —
rainwater ............................. 24
11. 618O of sulfate in rainwater: calculated compared to measured .... 28
12. Percent of sulfate formed by peroxide oxidation, assuming negligible
primary sulfates ......................... 30
vn
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ACKNOWLEDGMENT
The authors wish to acknowledge that some of the laboratory studies involving the
determination of isotopic relationships using solutions of known 18O/16O ratios were con-
ducted under a related program for the U.S. Department of Energy's Office of Energy
Research, Basic Energy Sciences, Division of Materials Science. Also, the authors are
grateful to Dr. Mark Dubois, Whiteface Mountain Field Station, Atmospheric Sciences
Research Center, Wilmington, NY; to Dr. Gary Eaton, Research Triangle Institute, Re-
search Triangle Park, NC; and to Dr. William Pierson and Ms. Wanda Brachaczek, Ford
Research Engineering Staff, Dearborn, MI, for collection and shipment of rainwater samples
to Argonne for analysis.
vin
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SECTION 1
INTRODUCTION
Considerable interest has been shown in the possibility of using oxygen isotopy to
elucidate the role of H2O2 in the oxidation of SO2 to sulfates in the atmosphere. The
potential importance of II2O2 lies in the view that aqueous-phase oxidation of SO2 to
II2SO4- probably accounts for a major fraction of the observed SO^" in the precipitation
occurring in the northeastern United States. The key reactants responsible for this ox-
idation are not well known, although a large number of possible catalysts and oxidants
exist in the atmosphere, including carbon, transition rnetal ions, hydroxyl and organic free
radicals, hydrogen peroxide (including organic peroxides) and ozone (Penkett et a/., 1979).
It has been postulated that the atmospheric oxidation of SO2 is limited by the availability
of these oxidants and catalysts, rather than by the availability of SO2 itself. This is the
so-called phenomenon of non-linearity between the concentration (and, by implication, the
emission) of SO2 in the atmosphere and the conversion of this SO2 to H2SO4. If such is
indeed the case, it is important to determine if peroxides have a pivotal role in acid for-
mation and deposition from the atmosphere [oxidation by ozone decreases in significance
with decreasing pH of the rainwater].
Oxygen isotopic studies are useful in distinguishing between the different oxidation
mechanisms effective in the atmosphere. Laboratory simulation of several different atmo-
spheric reaction sequences has shown that the oxygen isotope ratio in the product SO^"
is uniquely related to the reaction pathway followed in its formation (Holt et al., 1982).
It may therefore be possible to determine if the atmospheric hydrogen peroxide is respon-
sible for significant oxidation of SO2 to sulfate, and if it is this oxidant that limits the
aqueous-phase formation of sulfuric acid. The results of these studies could have signifi-
cant implications for energy technology, particularly if they indicate that, because of the
non-linearity in SO2 conversion discussed above, it may be more important io reduce the
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ambient concentrations of peroxides than of SO2.
In the 1981 JASON Committee Report to the U.S. Department of Energy (Chamberlain
et al., 1981), some recommendations for further research were based, at least in part, on
our earlier work on SO2 oxidation by H2O2 (Holt et al., 1981). According to our earlier
work, the 518O [deviation in parts per thousand (%o) of the !8O/lflO ratio of the sample
from that of the standard reference material, Standard Mean Ocean Water (SMOW)] of
sulfates produced by H2O2 oxidation were significantly lower than the <5I8O of sulfates
found in rainwater. However, the 518O of the reagent-grade H2O2 used in those exper-
iments was not known. The results suggested the need for isotopic analysis of H2O3 in
dilute solutions, and for a methodology whereby the <518O values of H2O2, H2O, and SO^~
in rainwater could be compared, in order to assess the importance of H2O2 in the formation
of sulfate-constituted acid rain.
The plan of this investigation was to develop a method for the determination of the
<518O of H2O3 in dilute aqueous solutions (simulating rainwater); to prepare solutions of
H2O2 of various 18O enrichments; to use the freshly prepared solutions of H2O2 to oxidize
SO2 to SO^" for evaluation of the relationship between 618O~,-,2_ and 618Opj Q ; and to
apply this relationship to the measured 618Ojj Q , 518Oo/-wa-, and £18Ojj Q in precipi-
tation water, for assessment of the importance of H2O2 in the atmospheric transformation
of SO2 to sulfate.
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SECTION 2
EXPERIMENTAL
Development of Analytical Method
No suitable analytical method was available for the quantitative extraction of the
oxygen in H2O2, (dissolved in water, ppb range), for isotopic analysis. Figure I shows a
4-step method which we developed for this purpose. It consists of the removal of dissolved
O2 from 20-liter samples of water by a combination of evacuation, ultrasonic agitation, and
sparging with helium; oxidation of the dissolved H2O2 to O2 in the water with KMnOv,
removal of the newly formed O2 from the water in a carrier-gas stream; and conversion of
the O2 to CO2 by reaction with platinum-catalyzed carbon at 600°C. The CO2 is then
mass spectrometrically analyzed to determine its 618O (identical to that of the oxygen in
the original H2Oi).
Analytical Train. Diagrams of the apparatus are shown in Figs. 2 and 3. Fig. 2
shows the all-glass analytical train. It consists of a bed of molecular sieve at —196°C for
the removal of traces of O2 from helium carrier gas, a gas pipet for the injection of known
amounts of O2 or CC>2 into the carrier gas stream during standardization, the water sample
chamber detailed in Fig. 3, a 20-mm o.d. cold trap at —78°C for the removal of residual
water vapor from the gas stream, a cold trap at -196°C for the removal of CO2 that is
scrubbed from the water sample by the helium stream, a bed of activated charcoal (3 g,
8-10 mesh, coconut grade, in a vertically mounted quartz tube, lined with platinum gauze)
at 600°C for the conversion of O2 to CO2 in the carrier gas stream, another cold trap at
— 196°C for the collection of the newly formed CO? from the gas stream, a capillary open-
well mercury manometer for measurement of the CO2, and a gas sample bulb, attached to
the train for transfer of the CO2 to a mass spectrometer for isotopic analysis. A cold trap
at —19G°C, not shown in the diagram, is in (he vacuum manifold to protect the analytical
train from vapor contamination from the mechanical vacuum pump.
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SAMPLE
H?0
Dissolved 02
H202
Dissolved 02
OXYGEN
REMOVAL
H2
H2
- Low prssaura H* spnrg«
•f utiraionic ogtlation
co2
To Mass Spec
Analysis
02 -» C02
Chorcool/Pt b«d
600°C
KMnO,
H202
OXIDATION
H20
OXYGEN
RECOVERY
Low prassurt H« iporg*
•+• ultrasonic agitation
Figure 1. Four-step method for conversion of oxygen in
(ppb range in water) to COa for mass spectrometric
analysis
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REACTION
CHAMBER
FLOW
METER
HELIUM
MOLECULAR
SIEVE
(-I96-C)
COLO
TRAP
(-I96-C)
COLO GAS
TRAP SAMPLE
(-!96-C5 BULB
CHARCOAL
PLATINUM
FURNACE
(600'C)
COLO TRAP
(-78-C)
CAPILLARY
MANOMETER
COLO TRAP
(-I96-C)
Figure 2. Analytical train
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ANALYTICAL
TRAIN
COLO WATER
CONDENSER
KMn04
CRYSTALS
20-LITER FLASK
ULTRASONIC BATH
Figure 3. Reaction chamber
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Reaction Chamber. The reaction chamber assembly, Fig. 3, consists of a 20-L round
bottom flask, a rotatable side-arm tube for the addition of crystalline KMnO4, a cold
water condenser to limit the amount of water vapor swept by the helium stream into the
—78°C cold trap, three stopcocks for manipulation of the helium stream, and a funnel
for the addition of solutions of acid and oxidants (used only during the development of
the method). The round-bottom flask was supported by a stainless sleel rack in an 83-L
stainlfss-stcel tank, interior 40.6 x 50.8 x 40.6 cm, (Model ATI! 1620-21), of an ultrasonic
cleaning system (Model EMa 50-24) manufactured by Branson Cleaning Equipment Co.,
Shelton, CT. The water in the ultrasonic tank was maintained at about 10°C by circulation
through a refrigeration unit.
The components of the glass analytical train were connected by No. 18 ball joints,
sealed with solidified black wax. The hollow plug of stopcock 22 was modified to provide
extra volume in the capillary manometer (Holt, 1955).
Conversion of O2 to CO2. The first part of the method to be developed was the con-
version of O2 to CO2 in the carrier gas stream (4th box in the flow sheet, Fig. 1). The
procedure is a modification of an established vacuum technique for the conversion of oxy-
gen in air to CO2 for mass spectrometric analysis. By the vacuum procedure, the air is
circulated at low pressure over graphite and platinum at 600°C and the resulting CO2 is
cryogenically removed from the closed system (Horibe et al., 1973; Holt, 1977). By our
new method, in which helium is conducted through a bed of 8-10 mesh activated coconut
charcoal at 600°C and in contact with platinum gauze, the yield of CO2 was found to be
affected by both the bed temperature and the axial temperature gradient. With a suitably
long furnace to obtain an essentially isothermal bed, the recovery of O2 (as CO2) as a
function of bed temperature is shown in Fig. 4. The data show that the optimum bed
temperature for maximum yield of CO2 was ~600°C, the same as that used by Horibe
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(SI
o
o
100-
90-
80-
70-
(f>
o
CM 60
O
CD
>
O
o
50-
40-
C£ 30-
20-
10-
0
0 100 200 300 400 500 600 700 800
Temperature of CC-Pt (°C)
Figure 4. Recovery of Oj as COa after reaction with charcoal
in contact with platinum at various temperatures
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et al., (1973) and Holt (1977) in vacuum-line applications.
The data in Fig. 4 also indicate that the maximum recovery of the added O2, measured
as CO2, was ~Q7%. Recoveries higher than 97% were apparently unattainable because
of competing reactions. At temperatures lower than 600°C, the formation of CO2 was
decreased due to the incomplete reaction,
02 + C->C02 (1)
and at higher temperatures, the yield of CO2 was decreased by the competing reaction,
+ C-»2CO. (2)
The effect of reaction (2) was demonstrated by injecting CO2 into the carrier gas instead
of O2. The results, Fig. 5, confirm that about 3% of the CO2 decomposed at 600°C.
The 97cc recovery is adequate for isotopic studies; and, at 600°C, oxygen isotope
fractionation within the equilibrated system of CO2, CO, unconverted CO2, and fixed
oxygen on the charcoal is negligible. The demonstrated reliability of the 979o recovery at
600°C allows the applicability of the technique to the quantitative determination of O2 in
gas streams, as well as to isotopic studies.
Removal of Dissolved Oq from Water
After it vas experimentally demonstrated that O2 could be reliably converted to CO2
in a carrier-gas stream, the technique for the quantitative removal of dissolved oxygen
from multi-liter quantities of water was developed (1st and 3rd boxes in Fig. 1). Various
combinations of sparging with helium, vacuum pumping, and ultrasonic agitation were
tested. The best results were obtained by a combination of all three. (The commonly used
degassing method of alternately freezing and thawing the water under vacuum was much
too impractical and time-consuming for large samples of water.)
-------
o
O
O
100
90H
80
70-
O
O
50-
40-
30-
20-
10-1
0
v-
-V-
v.
A First run after increase of temperature
V First run after decrease of temperature
0 100 200 300 400 500 600 700 800
Temperature of CC-Pt (°C)
Figure 5. Recovery of COj after exposure to charcoal and
platinum at various temperatures
10
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Helium Sparge. Attempts were made at first to remove dissolved oxygen from water
(saturated with air) by only sparging with helium. Quantitative removal of air oxygen from
the water by this technique was intolerably slow. Curve SP-1 in Fig. 6 shows the percent
removal of O2 (estimated to be dissolved in the air-saturated water at room temperature)
plotted vs. minutes of sparging. After 50 min of sparging, only ~l/3 of the dissolved O2
had been removed.
T. G. Holt (1983) cited a comparison of various degassing techniques in Technical
Brief No. 101 of Waters Associates, Milford, Massachusetts. The techniques, listed in the
order of decreasing effectiveness for the removal of dissolved gases from liquids, were: (1)
vacuum pumping combined with ultrasonic agitation, (2) vacuum filtration using a Waters
Solvent Clarification Kit, (3) ultrasonic agitation, (4) vacuum pumping, (5) boiling, and
(6) sparging.
Vacuum-Ultrasonic Agitation (VU). The apparatus was arranged to accommodate
treatment of the water sample by a combination of vacuum pumping and ultrasonic agi-
tation. The results of two experiments, VU-1 and VU-2, are plotted in Fig. 6. The rate of
removal of dissolved oxygen was found to be greatly enhanced by "rinsing" the water with
helium. That is, the evacuated sample chamber was filled with pure helium to atmospheric
pressure by allowing the gas to bubble through the water; then it was exhausted through
the analytical train. The sequential dissolution and removal of the helium from the water
had the effect of "rinsing" the O2 from the water. Curve VUR-1 in Fig. 6 shows the
improved results of a third experiment in which the water sample was rinsed with He at
regular intervals throughout the O2-removal treatment.
Vacuum-Ultrasonication-Sparging (VUS). The improved rate of O2 removal by fre-
quently rinsing the O2 from the water with helium suggested that a low-pressure sparge
by helium should be combined with the vacuum-ultrasonic agitation treatment. This tech-
nique, tested in runs VUS-1, VUS-2, VUS-3, and VUS-4 yielded the best O2 removal rates.
11
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X
O
CD
~0
to
O
«*—
O
"o
o
E
CD
100-
80-
60-
40-
20-
vus-i,?
S = Sporgtng
VU
•^ Ul trosonlcal Ion
VUR B VociA^>
+ UlIrasoolcatlon
+ HjIiun "rInso"
= VQCULITI
4- Ul tr aiootcat Ion
+ H« I iun Sparging
40 60 80 100 120 140
Degassing Treatment Time (min)
160
180
Figure G. Removal of O3 from water by various degassing
techniques
12
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The results show removal of ~99% of the dissolved oxygen in 1 h of VUS for 3-L water
samples. (20-liter samples were later found to require ~5 h of the treatment for complete
air removal.)
Oxidation cf H2O2 to O2 in Water
Bromine water was tested as an oxidant of H2O2 to O2 in the degassed water sample.
Although it was very effective in converting a known amount of H2O2 to O2, its use was
complicated by its exposure to stopcock grease and by contamination of the analytical
train. Curve "VU-1 + H2O2 + Br2" in Fig. 6 shows the results of cine experiment in
which a degassed water sample (VU-1) was spiked with H2O2 and treated with bromine
water, yielding a recovery of about 97% in 50 min of VU treatment.
Potassium permanganate was used in all subsequent experiments. Curve "VUS-2 +
H2O2 + KMnO4" in Fig. 6 shows the results of an experiment made to compare its
effectiveness to that of bromine. The recovery of oxygen, added as HjO2 to a degassed
water sample (see curve VUS-2), was also ~979& after 1 h of vacuum ultrasonic sparging.
In earlier experiments with KMnC-4, it was added as a O.lN solution through the funnel,
Fig. 3; later, the rotatable side arm was made a part of the apparatus to permit the
addition of KMnC"4 crystals, with correspondingly lower procedure blanks.
Synthesis of 18O-enriched H2O2
Hydrogen peroxides of various <518O were not commercially available. A suitable
method of synthesis (Vol'nov et al., 1964) was identified and successfully applied to the
laboratory preparations of four stock solutions of H2O2 of different 618O. By this method,
H2O2 is formed by exposure of supplies of water vapor, each differing in 518O, to a high-
voltage (~1.4 kV) discharge in ~100 cm of 10-mm o.d. glass tubing between two water-
13
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cooled aluminum electrodes. Some of the HO radicals, formed by the dissociative reaction,
2H20 «=i 2110 + 2H (3)
are condensed in a cold trap (-196°C) where they combine to form H2O3, leaving the H
radicals to combine in formation of H2 and be pumped away through the vacuum line.
Other techniques which we experimentally found to give inadequate yields of HaO2
were conduction of an electric arc across a stream of aerosolized water droplets (Kok,
1982), excitation of water vapor by a radio-frequency silent discharge in a glass chamber
(7.5 cm dia x 20 cm long) in a commercially available plasma cleaner unit, and excitation
by a glow-discharge unit (4.8 cm dia x 70 cm long) that had uncooled aluminum disk
electrodes (Jarnagin and Wang, 1958).
Collection of Rain Samples
Each rain sampling station consisted of four 1-m2 plastic skylights, inverted to form
funnels and fitted with plastic nipples and hoses to allow the water to flow into two 25-L
plastic collection bottles. From each collected sample of >29-L, a 4-L bottle was filled for
subsequent determination of 618Opj Q anc' ^18O^-/~,2-, and a 25-L bottle was filled for
subsequent determination of <518Oo Q . The 4-L sample was treated with 20 mg CuCl to
prevent bacterial alteration of the sulfate during storage before analysis; the 25-L sample
was treated with 25 ml concentrated H2SO4 and was refrigerated at ~5°C to prevent
autodecomposition of the H2O2 during storage before analysis.
During the period June-August 1985, rainwater collection stations, in addition to
the one at Argonne, IL, were operated by the Research Engineering Staff, Ford Motor
Company, Dearborn, MI; Research Triangle Institute, Research Triangle Park, NC; and
Whitcface Mountain Field Station, State University of New York, Albany, NY. Approxi-
mately two rain evt-iits per month were sampled at each site during the 3-month period.
14
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The rain collectors were installed by GP Engineering, Downers Grove, EL, at each of
the three sites. Insulated containers were used for shipment of 25-L bottles of acidified,
chilled rainwater by air express to Argonne for analyses.
15
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SECTION 3
RESULTS
Recovery of added H2O2
The recovery of reagent grade H:O2) added to 20 liters of distilled water, was about
96%. This recovery is adequate for determining the oxygen isotope ratio and the concen-
tration of HoOa in rainwater.
Test for Isotopic Interference
Schumb et al. (1955) cite references showing that dissolved HaO2 does not exchange
oxygen atoms with the solvent water, dissolved molecular oxygen, or oxygenated products.
In decomposition of H2O2 by oxidation to molecular oxygen, the O-O bond is not broken
and no fractionation occurs. The MnO^" ion does not cause isotopic interference (Cahill
and Taube, 1952). However, in decomposition of H2O^ by reduction (not applicable to our
experimentation), the bond is severed and fractionation may occur.
To confirm the absence of appreciable isotopic interference in our procedure by oxygen
exchange between the H2O and either the H2O2 or the O2, before, during, or after the
oxidation reactions, the oxidation was carried out in the presence of three different water
supplies of various 518O. The results in Fig. 7 show that the 619O of the CO2 product
was unaffected by the 518O of the water solvent.
Effects of Autodecomposition of H2O2
Results from two sets of experiments on the autodecomoosition of H2O2 (spikes added
to 20-L samples of rainwater), over storage times of up to 11 days, showed that if the
rainwater is stored unacidified at room temperature, the H2O2 concentration declines
rapidly (~30% depletion in 2 days and ~9S% in 11 days); and that the <518O of the
16
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0
-------
undecomposed H2O2 undergoes a corresponding increase (~5%, during the first 5 days).
However, if the rainwater is first acidified (~20 ml cone. H3SO4 in 20-L) and kept cold
(1-9°C) during storage, the changes in the concentration and in the <518O of the H2O2 are
negligible.
Effects of KMnO4 Decomposition
Decomposition of excess KMnO4 by organic matter in rainwater was shown to product-
only CO2, which is cryogenically removed in the analytical train before the H3O3-derived
O2 is converted to CO2. The reaction between KMnO4 and organic matter does not
produce O2 and is th°refore not a source of interference in the method.
Procedure Blank
The blank of the analytical procedure was reduced from 8.5 to 3.5 /imoles of O2 by
adding the KMnO4 to the 20-L sample as pulverized crystals rather than as a pre-boilod
KMnO4 solution. (In 20-L of water, 3.5 /^moles is ~3 ppb.)
Isotopic Relationships
High-Voltage Preparation of H3O2 from H2O. The dependence of the 518O of M3O3
on the 518O of the H2O from which it was prepared by the high-vol*age discharge method
described above is shown in Fig. 8. The equation of the best fit regression curve is
*18°H203 = 1.03<5'80H20 +29.4%,. (4)
The results indicate that the <5'8Ojj2Q3 is controlled directly by the <518O of the H3O
and that the 518O of the H2O2 is substantially higher (~29%o) than that of the water
from which it was formed. Although the high-voltage process used in these preparations
18
-------
60
50-
CM 4°-
O
CM
^
*s 30:
o
00
40 20-
10
0
0
10
-20 -10
1R
6 0 of H20 (precursor)
H
20
Figure 8. Dependence of 518Oj|3Q,j on S16O^Q in a high-
voltage discharge process ol formation
19
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of H2O2 may simulate lightning in the atmosphere, the importance of lightning, relative
to other sources of H2O2 in the atmosphere, is open to question (Kok, 1982).
H2O2 Oxidation of SO2 to SO^". Using the four stock solutions of hydrogen peroxide,
each of different 618O, sulfate solutions of correspondingly different <518O were prepared by
oxidation of SO2 which was in isotopic equilibrium with water of constant <518O = —7.9%,.
In Fig. 9 the 618O of each resulting sulfate is plotted versus the 618O of the H2O2 and the
equation of the best-fit regression curve is
5l8°SOr = °'43 518°H203 + 3-5V (5)
The regression curve of the previously determined relationship (Holt et al., 1981)
between <518Oc/~)2- and <518Oj^ Q in aqueous-phase oxidation of SO3 by H2O3 was
*- = 0-57 5180H20- 2.4%,. (6)
Assuming that all significant effects of the <518O of the SO2 on the 618O of the SO^~
are lost by rapid isotopic exchange between the SO2 and the large excess of water, prior
to appreciable oxidation (Holt et al., 1981), 618Ojj Q an(^ ^'8^H O rema'n as ^e on'y
two complementary variables in the equation for <518OQn2-; therefore, the comprehensive
^»J4
regression curve for <518Ocr.2_ is
)2_ = 0.57 5180H20 + 0.43 5i80Ha02 + C (7)
The constant, C, was evaluated at 8.4%, from the data given in Fig. 9 by substituting the
corresponding measured values for (518OgQ2- and <$18OH2O2' and ~7-9%> for <5l8°H2O-
The comprehensive equation then becomes
6180S()2_ = 0.57 <5180H30 + 0.43 5180H202+ 8.4%, (9)
20
-------
30
•20-
CM -
-------
and can be used to calculate 61BO^^- from 518Ojj Q and £18Opj Q of rainwater for
comparison with corresponding measured values of $18Oq~a_. This comparison may prove
to be uniquely useful in the assessment of the importance of HaO2 in the oxidation of SO2
to SO^" in the atmosphere.
The slope of 0.43 (approximately 2/5) in Eq. (5) of the regression curve through the
data of Fig. 9 confirms the evidence of the intermediate species, I^CvSOg", which was
previously proposed (Holt et a/., 1981). Apparently, the 618O of the sulfate product is
2/5-controlled by the two oxygens in the H2O2 of the adduct, and 3/5-controlled by the
HSO^~, which, in turn, is isotopically controlled by rapid oxygen exchange with the large
excess of water with which it is associated.
6IBO and Concentration of II2O2 in Rainwater
Forty-four samples of rainwater, collected at Argonne from September 1984 through
November 1985 and at the other three sites (in Michigan, New York, and North Carolina)
during the summer months of 1985, were analyzed by the new method. Results obtained
for 14 of these samples arc given in Table 1 and in Fig. 10. Isotopic data for H2O, SO^",
and Il2O2 in rainwater from the four sites are plotted vs. time in Fig. 10; in addition, Table
1 gives concentrations of H2O2 and SO^" in the rainwater, and the elapsed time between
collection and analysis of each sample. Our analytical method d >es not discriminate H2O2
from organic peroxides in rainwater; consequently, in the report of field results and in the
discussions that follow, "H2O2" will refer to combined H2O2 and other active peroxides.
Results for 25 of the samples are not reported because the concentrations of H2O2 were
too low (<15 ppb) to yield sufficient amounts of CO2 to give reliable isotopic data by mass
spectrometric analysis. Discarding of these samples of low H2O2 concentration has the
effect of estimating a lower than actual proportion of sulfate formed by peroxide oxidation
of SO2 discussed later in the report. [The cause(s) of very low peroxide concentrations in
rainwater samples may be one or more of the following: (1) during early use of the method,
22
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Table I.
Concentration and isotopic results for rainwater collected at Argonne
National Laboratory, IL, (ANL); Dearborn, MI, (DBA); Whiteface
Mountain, NY, (WFM); and Research Triangle Paik, NC, (RTF);
May-September, 1985.
. Concentration 518O
Sample No.
ANL-5/27/85
WFM-6/07/85
RTP-6/ 12/85
RTP-6/ 18/85
ANL-7/02/85
DEA-7/02/85
ANL-7/09/85
RTP -7/10/85
WFM-7/ 15/85
RTP-7/ 16/85
WFM-7/31/85
DEA-8/ 14/85
RTP-8/20/85
ANL-9/06/85
Flotage
(days)
2
6
6
9
7
9
7
8
17
20
64
127
79
76
H202
(Ppb)
330
124
147
374
316
23
90
433
201
224
160
16
184
33
so*-
(mgL-1)
2.4
1.5
2.4
2.1
2.1
4.4
4.4
3.6
2.0
3.6
1.2
2.9
4.1
1.2
H202
44.8
50.5
47.5
45.0
41.4
58.3
54.2
53.8
48.4
58.0
41.0
47.9
49.6
44.8
H20
-3.1
-3.7
-3.5
-2.5
-3.1
-6.6
-1.0
-2.6
-3.9
-1.5
-7.9
-1.3
-3.5
+0.4
^ V^y >i
15.8
15.0
14.4
15.8
12.6
12.9
16.9
14.3
15.3
16.4
14.1
16.8
12.8
16.1
23
-------
o
o
O
00
60-
50-
40-1
30-
20-
10-
0-
-10
O Argonne A Whiteface Mountain
D Dearborn V Research Triangle Park
so:
H20 O
1985
Figure 10. Measured isotopic data for H2C>2, SO4 , and H2O
in the samples of rainwater
24
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storage procedures were inadequate for prevention of autodecomposition of peroxides; (2)
several samples were collected during winter months of 1984-85 when concentrations of
peroxide? in precipitation were very low, and (3) high concentrations of SOj, relative to
peroxides, in the atmosphere depletes the peroxide that might otherwise be present in
collected samples of rainwater.] Similarly, three samples were disqualified because of very
low sulfate concentrations (<0.3 mg L—1), and the data for two other samples were not
included because of analytical difficulties (back diffusion of CO2 from the charcoal furnace
into cold trap 2, when the line was inadvertently evacuated between the furnace and the
water sample with no helium flow).
25
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SECTION 4
DISCUSSION
»518O of H2O2 in Rainwater
A unique characteristic of atmospheric peroxides, as shown for the first time by ap-
plication of our newly developed analytical method, is the very high <518O, relative to the
618O of atmospheric water, sulfates, and air. As demonstrated in Fig. 10, the measured
<518O of atmospheric peroxides ranged from 45 to 60%,; sulfates, IS to 17%,; and rainwater
0 to — G%O. (The <518O of air oxygen is constant at 23.5%0.) In contrast to the very high
<518O of atmospheric peroxides, the <518O of reagent-bottle H3O3 (Fisher Scientific, H325)
was found to be —6%,.
The oxygen isotopy of peroxide formation in the atmosphere, resulting in such high
S1SO, is intriguing because it is necessarily related to the mechanism(s) by which the per-
oxide is formed. For the various sources of atmospheric peroxides (Lee, 1985), each possible
mechanism of formation is characterized by its own oxygen isotopy. Consequently, our re-
sults suggest that oxygen isotopic studies may well be uniquely applicable to investigations
of the origin of peroxides in the atmosphere.
As mentioned earlier, the 618O of peroxides in rainwater was higher than might be
expected for H2O2 that is formed by high-voltage discharge, such as lightning. For example,
application of Eq. 4 to the range of 618O in water (Fig. 10) would lead to an estimated
range for H3Oa of ~24%> instead of the observed 45-60%,.
A phenomenon that may contribute substantially to the relatively high <518O of per-
oxides in rainwater (depending on the conditions of the rainfall, the sample collection, and
the storage) is autodecomposition of the peroxide. As the peroxides decompose (under
favorable conditions of catalyzing contaminates, temperature and pH), the <518O by the
residual peroxides is expected to increase. Therefore, depending on the extent of autode-
composition between the time of liquid-phase oxidation of SOa in the atmosphere and the
26
-------
time of collection, acidification, and refrigeration on the ground, the 61BO of the residual
H2O2 in the analyzed sample may be significantly higher than that of the H3O3 involved
in the SO2 oxidation. Further experimentation to establish the relationship between the
618O of peroxide in cloud water and that of concurrently collected rainwater on the ground
might provide a correction factor which would improve the significance of the isotope data
that is conveniently available from the analysis of water in ground-based collectors.
618O of Atmospheric Sulfates: Measured vs. Calculated
The measured 618O values for sulfates, Fig. 10, are co-plotted in Fig. 11 with corre-
spondingly calculated values for HjC^ oxidation, metal-catalyzed aqueous oxidation, and
primary sulfates. The measured values are clearly higher than those calculated for metal-
catalyzed oxidation and clearly lower than those calculated cither for H2O2 oxidation or
for primary sulfates. Calculated values for H2O2 oxidation were obtained by use of Eq. (9)
for metal-catalyzed aqueous oxidation by
6180so2_ = 0.86180H20 + 10%, (Holt et al., 1982) (10)
and for primary sulfates by
<518OgO2_ = 0.015<518OH2o +45%, (Holt et d., 1984), (11)
In 1982, Holt et al proposed that since their measured values for 618Oc~2- in precip-
itation water ranged consistently between the calculated values for primary sul.'-tes and
secondary sulfate? produced by metal-catalyzed oxidations, the fraction of atmospheric
sulfates at a given site originating as primary sulfates, could be estimated. At that time,
however, the isotopic qualities of atmospheric peroxides and of the sulfates which they
might produce were unknown. Since then field experiments were performed near a strong
source of primary sulfates, the isotopic results of which indicated that scavengement was
~300 times more efficient for sulfates than for SO2 (Holt et d., 1983). Consequently, we
expect that, during precipitation, scavengement of primary sulfates is essentially complete
27
-------
I
-------
within a few kilometers of the source, whereas beyond a few kilometers, the scavenged
sulfates are mainly secondary sulfates, formed earlier or within the storm system by one
or more mechanisms of SO2 oxidation.
Fraction of Atmospheric Sulfates Formed by Peroxide Oxidation
Assuming that primary sulfates of characteristically high 618O (Holt et al., 1984) are
effectively scavenged by rain within a short distance from their sources, an estimation of the
fraction of sulfates in rainwater several kilometers from a strong source can be made from
relative deviations of the measured <518Oc;r.2- values from the corresponding calculated
oU4
<518O<-,,^2- values for peroxide oxidation and metal-catalyzed O2 oxidation, respectively.
Using the following relationship for calculation of the percentage of sulfate in rainwater
formed by H2O2 oxidation,
4 (meaa.) 4 (calc., O2 oxdn.)
100'
4 (calc., H2O2 oxdn.) b(J4 (calc., O2 oxdn.)
the percentages for our four sampling sites during 1985 are plotted vs. time in Fig. 12.
The average for all of the values at all of the sites is 37 ± 8% sulfate, formed by peroxide
oxidation.
Equation (12) assumes that all sulfate measured in the rainwater samples was pro-
duced within the raining cloud and that only H2O2 and metal-catalyzed aqueous oxidation
contributed to the oxid if ion of the dissolved SO2. Other potential sources of sulfate in the
rainwater Camples were aqueous-phase ozone oxidation of SO2, dissolution after gas-phase
reaction of SO2 with OH radicals, and dissolution of sulfate nuclei previously formed in
non-precipitating clouds by any of the above oxidation processes.
In reviewing our work, Dodge (1986) pointed out that the results of other researchers
(Scott, 1982; Hegg and Hobbs, 1984; Scire and Venkatram, 1985) variously indicate that 65
to 85% of sulfate in oloudwater is due to in situ aqueous phase reactions and the remainder
29
-------
100
90-
80-
70
60-
•z.
UJ
O
C£
UJ
Q_
40-
30
20-
10-
O Argonne A Whiteface Mountain
D Dearborn V Research Triangle Park
O
V _ „ , -A
~
•-- o
1985
Figure 12. Percent of sulfate formed by peroxide oxidation,
assuming negligible primary sulfates
30
-------
to cloud scavenging of pre-existing sulfates. However, since the residence time of aerosol
sulfate in the atmosphere is rather short, any aerosol sulfate scavenged could be expected
to be about the same in isotopic quality as that in the precipitating cloud system. Further,
ozone is expected to play a minor role in the oxidation of SO2 in cloudwater of pll lower
than about 5 (Penkett et a/., 1970).
At best, these data establish a lower limit for peroxide oxidation. As discussed above,
the <518O of atmospheric peroxides may be somewhat lower than that which is finally
measured from the ground-based rain collector. It follows that the calculated 618O^~2-
would be correspondingly lower, and the *>o of sulfate formed by peroxide oxidation would
be correspondingly higher.
The technological significance of our findings (that ~40% or more of acid sulfate
in rainwater originates from peroxide oxidation of SOa) is that further investigations of
the origin(s) and of possible methods of control of atmospheric peroxides are of prime
importance.
31
-------
SECTION 5
SUMMARY
A new method was developed for the determination of the 618O of peroxide (ppb
range) in rainwater. Experimental results showed the method to be reliable with respect
to recovery of added HaO^, blanks, and freedom from isotopic interference. A procedure
for collection and treatment of 25-L samples of rainwater with minimum peroxide de-
composition between collection and analysis was developed. An isotopic relationship was
established between the oxygen in sulfate and the oxygen in the peroxide and in the water
involved in the peroxide oxidation of SO2 to SO^". By using this relationship for peroxide
oxidation and a similar relationship previously determined for aqueous metal-catalyzed
oxidation, the fraction of sulfates formed by peroxide oxidation ;n rainwater from four
sampling sites in Illinois, North Carolina, New York, and Michigan was estimated to be
~\Q%. This fraction could be substantially higher if a substantial amount of peroxide
undergoes decomposition between the time of oxidation of SOa in the atmosphere and
the time of collection and peroxide-stabilization treatment of the rainwater sample on the
ground. Our results show that atmospheric peroxides play a major role in the formation
of sulfates in the atmosphere, and, therefore, that it is essential to obtain all possible
information on their origin and control.
32
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