by invoking a charge balance equation. The formation of sulfate
can then be calculated by a simple integration of the assumed
first-order oxidation rate equation

a[so,]

ae = * (S031.

1 1

The rate constant k was taken to be 0.1 minTM~ or 0.0017 sec”
based on an analysis of van den Heuvel and Mason's data. The
calculations made under these assumptions did not show the sul-
fate formation to reach an asymptotic limit in one to two hours
as noted by Junge and Ryan, nor did it show the linear depend-
ence on the initial partial pressure of SO,. The work yields
oxidation rates of the order of 2.5 percent per hour and does
suggest sulfate levels of the same order of magnitude as some

7), but the corresponding NEY concentrations and

measurements
pH's are much higher than measured, and in general rather long
times are required for large sulfate formation. Depletion of

the gas phase concentrations of SO, and NH was not considered.

Assuming the same set of reactions and equilibrium constants as

Scott and Hobbs, Mckay (58)

was able to predict much faster for-
mation of sulfate by calculating the first order rate constant
from the expression given by Fuller and Crist:

k = 0.013 + 59(H*]1/? sec? .

At pH 7, this is 20 times as great as the value used by Scott

and Hobbs, and at pH 5, it is 120 times as great. Furthermore,

56


three-year program to measure the degree of conversion of SO,
and to evaluate its impact on atmospheric sulfate concentrations
and visibility. So, conversion rates were obtained by determin-
ing the difference between upWind and downwind concentrations of
particulate ammonium sulfate relative to ambient So, and wind
speed. Measurements in these tests were made 13 and 37 miles
downwind from the plant. Conversions of SO to particulate

sulfate were as follows:

Test 1 0.37%/hr 13 miles downwind RH 36-51 percent
Test 2 0.76%/hr 13 miles downwind RH 26-40 percent
0.45%/hr 37 miles downwind.

SO, concentrations ranged from 0.01-0.06 ppm in these tests.

These results are of interest for several reasons. They are
substantially lower than the oxidation rates reported by TVA,
possibly because of the low relative humidity at the Four
Corners location. Of greater significance, the resultant S05
oxidation rates are derived from direct, positive analysis of
sulfates and not inferred from the loss of SO, relative to
some intrinsic tracer in the plume. It is quite possible that
in this respect the Four Corners tests provide a more accurate
assessment of SO, oxidation rates in plumes.

A recent study by Davis and Klauber 22) | is cited here,

not so much for what it adds to our knowledge on SO,

23


EPA-450/3-76-022

SO, OXIDATION IN PLUMES:

A REVIEW AND ASSESSMENT

OF RELEVANT MECHANISTIC
AND RATE STUDIES

by
A Levy. DR. Drewes. and J.M. Hales
Battelle Pacific Northwest Laboratories

P.O. Box 99
Richland, Washington 99352

Contract No, 68-02-1982
Program Element No. 24C129

EPA Project Officer: Joseph A. Tikvart

Prepared for

ENVIRONMENTAL PROTECTION AGENCY
Office of Air and Waste Management
Office of Air Quality Planning and Standards
Research Triangle Park, North Carolina 27711

September 1976


and ky was found to be 3.32 + 0.11 x 10° wis}, This leads to

an oxidation rate of about 0.21%/min in a cloud under typical
conditions (10°, pH 5, about .1 tol g/m? liquid water content)
assuming 7 ppb S0,- The ozone level assumed, 50 ppb, represents
yet another extrapolation of the laboratory data. This rate is
some 70 times faster than that predicted by the analysis of
Scott and Hobbs, although it is of roughly the same order as
indicated by McKay's calculations, depending on the pH. Sub-
sequent work by Penkett and Garland (42) measured the rate of
oxidation of SO, by ozone in fogs formed in a chamber, and the

measured rates agree with those calculated in the earlier work.

It may be noted that in the above review, little attempt has
been made to reconcile the findings of various authors. This
stems partially from a feeling of frustration because there is
such a wide spectrum of results reported, and also because so
little of the work which has been done so far has been directly
concerned with the SO, in plumes. Even those studies which
attempt to relate their findings to the case of polluted atmos-
pheres are forced to make rather gross extrapolations from the
data collected in the laboratory, and the diversity of the re-
sults is hardly surprising. On the other hand, if there is one
fact which the literature does show well, it is that the oxida-
tion of S05 in the aqueous phase is highly sensitive to several

parameters, among which are pH, relative humidity, temperature,

65


the concentration of these ions on the surface, thus promoting

the conversion rate.

A study of the macroscopic properties of SO, adsorbed on mangan-

ese oxides '95)

is presented here mainly because of its relevance
to the catalytic role of Mn in aqueous systems discussed in a
previous section. This study looked at the adsorptive capacity
of wet and dry MnO, and Mn,03- Essentially, the study shows the
increased capacity for adsorption with increasing surface area
and decreasing specific gravity, both of which result from the
absorption of water.

(96)

A study by Happel and Hnatow is of interest here in relation

to some of the preceding studies. A theoretical analysis of the

18 35

use of the isotopic tracers 05 and SO in the catalytic oxi-

dation of SO over V205 catalyst suggests that in conversion
reactions in commercial acid plants and stack gas converters,
V205 is considerably more effective than in plumes. The study
concluded that oxygen chemisorption is the limiting mechanistic
step, with S03 desorption being of next importance in the rate
control. A rate equation is derived for conditions near equilib-
rium. It appears, however, that the rate equation is not applic-
able to ambient, possibly non-equilibrium conditions in a plume.

The analysis is useful, however, as a guide to investigators

applying isotopic tracers (e.g., Urone [90) and smith (91)),

75


SECTION IV

PLUME STUDIES

From a practical point of view, the most obvious or straightfor-
ward way to examine SO, oxidations would be to make measurements
directly in actual plumes. This approach has been carried out by
a number of laboratories, but unfortunately with minimal success.
Experimentally it is essentially impossible to control conditions
in the plume; added to this is the fact that such measurements are
very difficult to carry out. For this reason, results obtained
from such studies tend to conflict, making it difficult to draw
meaningful conclusions about plumes in general or to even compare
the studies to each other. Table 2 presents a summary of the

studies performed to date.

One of the first plume studies directed toward examination of

SO, - sulfate relationship, conducted by Gartrell et al., of the

) (12)

2
Tennessee Valley Authority (TVA , has become nearly classic,
in that investigators often attempt to relate theoretical and lab-
oratory experimental results to the TVA results. The TVA program
attempted to measure both SO, and SO, in the plume; however, the
instrumentation of that period (early 1960) was not well suited

*
to making So, {or H2S0,) measurements. Table 3, taken from the

*The inability to make simple, reliable measurements of ambient
H2SO, levels remains as possibly the most severe limitation to
studies of the fate of SO5 in plumes today.

13


In the Fe0, experiments there was rapid interaction with SO,
even at zero relative humidity. In the cases of Fe503 MnO, and
suspended particulate matter there was evidence for significant
physical adsorption of S05. In all cases, humidity increased the
adsorptivity of SO,.

Several species were inert to SO, i.e., did not adsorb S05;

these were Caco., V, Os and flyash. The V,0, results appear un-

3/ "2 2-5
usual since V205, is a recognized oxidation catalyst for SO,- At
500°C and above, it is presumed to oxidize via the reaction:
>
So, + V20 + SO, + V20,-

2 5
At room temperature, however, the sorption of SO, was only 1.2%

at 95 percent relative humidity.

A number of other studies have also been conducted on the inter-
action of SO, with the iron oxides Fe,0, and Fe,0,. In general,
these studies have limited applicability to resolving the prob-
lem of reactions in plumes in that rate information is limited.

Chun and Quon (89)

carried out a study designed to measure the
capacity of ferric oxide particles to oxidize SO, in air. The
procedure used involved laying down a film of Fe,03 on a filter
(by combustion of iron carbonyl vapor) and passing a stream of

SO, in air through the filter-reactor. The authors describe the

2
reaction as a "capacity-limited heterogeneous reaction", because

the reaction is not catalytic in the true sense of the term.

71


SECTION V

HOMOGENEOUS STUDIES IN THE GAS PHASE

Until rather recently it was generally felt that S05 must be
oxidized by heterogeneous (catalytic) paths and that the contri-
bution of homogeneous gas phase reactions was minimal at best.
Subsequently, however, there has been reason to challenge this
assumption. Although one cannot by any means downplay the
potential role of heterogeneous catalysis in So, oxidation,
several recent studies of SO, oxidation in power plant plumes,
cited in the previous section, show lower rates of oxidation
than were first reported --- rates well below two percent per
hour. It is quite possible that rates of SO, oxidation derived
from earlier studies were too high because of filter media ef-
fects and the poor analytical techniques available (problems
which persist to this day) as well as the indirect determina-
tion of sulfate formation. Thus, as the reported heterogeneous
rates get lower, the homogeneous reaction paths increase in
relative importance. This trend has approached the point, in
fact, that some investigators argue that homogeneous reaction
paths may be at least as significant as heterogeneous paths.

It is not our objective to argue this point here; the primary
point is that homogeneous reaction rates cannot be arbitrarily

assumed to be negligible compared to heterogeneous rates.

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limitations to such a device as a basic kinetic tool, a definite
influence of hydrocarbon concentration is being observed in these
studies. Rates of SO, oxidation., which appear to be homogeneous
and not especially influenced by the chamber walls, are of the

order of 2-6 percent per hour. Increasing the propylene from 1.6

ppm (as carbon) to about 3 ppm doubled the homogeneous rate.

Roberts and Friedlander (37) have recently carried out experi-
ments designed to study the formation of sulfur-containing
aerosols under ambient photochemical smog conditions. These
experiments were carried out in a large (96 m) Teflon chamber
irradiated with natural sunlight. - Seven olefins were used in
these studies, although most of the experiments dealt with l-
heptene-SO,-NO systems. Based on pseudo-first order deple-

tion of S05, the reaction rates varied from 0-90 percent per
hour, depending on the initial SO, level. The relative humidity
was quite low (< 40%), making the extremely high rates even more
surprising, although no explanation of the high rates was
offered. The authors do develop some theoretical kinetic
analyses using a mechanism similar to that of Cox and Peckett
(discussed earlier) and present an expression for estimating

aerosol sulfur content. The experimental results appear con-

sistent with a mechanism for the formation of the ozone-olefin

intermediate (I).

37


is insufficiently hydrated. The authors feel that a similar
mechanism may be involved where other catalysts are concerned,
although no others were used in either the experimental work or

theoretical development.

Further work on MnSO, and on MnCl5, Cuso,, and NaCl aerosols

1, (59),

was reported by Cheng et a In terms of catalytic effici-

ency, the following order was established: MnSO, > MnCl, > Cuso

4 4
> NaCl. In an attempt to extrapolate the experimental results to
atmospheric conditions, they assumed a fog of 15 um droplets con-
taining 0.2 g H,0/m> in which half of the droplets contain cata-
lyst at concentrations capable of oxidizing SO. at the same rate
as that found with 500 ug MnSO,/m°. Under these assumptions,

the extrapolation of their experimental ‘results indicated that

SO, at the 0.1 ppm level should be oxidized at a rate of about

2
20%/hr. The level of catalyst here was taken to be typical of
that in a power plant plume, and the resulting rate is about

four times higher than the .09%/min determined theoretically by

Foster.

A kinetic study of the homogeneous oxidation of SO, by cobalt
ions was performed by Chen and Barron (62), Although the value
for the rate constant was not determined, the reaction was
found to be zero order in oxygen, three halves order with re-
spect to sulfite ion concentration, and one half order with

respect to cobalt catalyst concentration. The mechanism proposed

61


*
also examined. Using fast-flow reaction techniques the pseudo-

bimolecular rate for homogeneous reaction was found to be 9.1 x

13 om? molec’ sec!; the actual termolecular reaction rate at

1.3 Torr and 300°K isthus 2.2 x 10779 om® molec * sect. (the

10

proportionality factor between the two values is, of course, [M]
at the experimental temperature and pressure.) Assuming a four-
center intermediate complex for the formation of H,S0,, this

appears to be a very rapid rate.

One of the basic questions in assessing the homogeneous rate of

SO oxidation in the atmosphere is how fast the reaction can pro-

ceed in a clean air system. Gerhard and Johnstone (27) | in their

early work in this area, reported an oxidation rate of 0.1 per-
cent per hour in natural sunlight and rates of the order of 0.68

°
percent per hour in U.V.-irradiated systems (2950-3650 A) with NOQ,-

(28)

Calvert's group considered the photooxidation reaction from

basic principles. Based on new estimates for the triplet so,

quenching rate with Noy Os, H,0, Ar, He and NO this group esti-

mates that the maximum rate for the homogeneous oxidation is 1.9

(29)

percent per hour. Friend disagrees with this estimate, believ-

*The intermediate of a reaction is a state in which the two react-
ants are bound in an energetic, short-lived complex. In many cases,
the intermediate is so highly energetic that it flies apart before
the rearrangements necessary for the formation of products can be
accomplished. However, if a third body (molecule or atom) collides
with the complex, it can draw off some of the energy, giving the
complex a longer lifetime and greatly increasing the chance of for-
mation of products; the reaction is then termolecular, and the third
body is designated M. In the atmosphere, of course, M is generally
N, or O

2 2°

33


others (52/79) |

The strong dependence of the rate in droplets on
relative humidity is seen to be due to the increase in pH (by

dilution) as the humidity increases.

Other Mechanisms

There have been few mechanisms proposed for the oxidation of So,
in aqueous systems other than the catalytic ones mentioned above.
Work on a photochemical process in the aqueous phase was reported

by Jones and Adelman (8?)

, but as is typical of such mechanisms,
the presence of reactive hydrocarbons is necessary. It is prob-
ably safe to say that photochemical processes are not important
in the aqueous phase oxidation of SO, in plumes under ordinary
circumstances due to the low concentrations of such hydrocarbons.
The possibility of entrainment of such species must be considered,
but kinetic studies of the aqueous phase photooxidation of SO.
are virtually nonexistent. Of possibly greater importance are
reactions with ozone (HSO,” +037 HSO,” + 0,). Espenson and

Taube (23

performed tracer experiments which, although they did

not yield rate data, indicated that the reaction is not as simple
as the stoichiometric relation above suggests, since oxygen atoms
are exchanged with the solvent as well as being transferred between

sulfite ions and ozone. penkett (61)

, using ozone in the range of
3 to 5 x 107° M, found the oxidation of bisulfite to be first
order with respect to ozone. The overall reaction was found to

fit a rate expression of the type - 4 (0,)/at = k, (05) (HSO,”)

64


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Here the K's denote equilibrium constants defined by the follow-

ing expressions:

Ko [S05]
P
SO,
- +
kK = {HsO, ][H']
1 ——_____
[S05]
= +
K = [S03 J(H"]

[HSO, ]

Values of the equilibrium constants have been measured experi-
mentally, and are given as functions of temperature in Table 9.
These values can be utilized in conjunction with the above
equations to demonstrate that, under typical environmental con-
ditions, the preponderance of sulfur dioxide dissolved in water

exists in the form of HSO,” ions.

The question of the rates of the dissociative processes (2) and
(3) has also been debated actively up until the very recent
past. Early semi-quantitative measurements indicating these
reactions to be extremely rapid were questioned in the sub-
sequent chemical engineering literature, and the quantitative

results of Wang and Himmelblau (49)

implied that the ionization
processes were sufficiently slow to be rate influencing under a

number of circumstances.

43


33.

34.

35.

36.

37.

38.

39.

40.

41.

42.

43.

44,

Smith, J. P., and P. Urone, Static Studies of Sulfur
Dioxide Reactions, Effects of NOo, C3H¢r and H20. Env.
Sci. Tech. 8: 742-746 (1974).

Wilson, W. E., Jr., and A. Levy, A Study of Sulfur Dioxide
in Photochemical Smog, I. Effect of SOj and Water Vapor
Concentration in the 1-Butene/NO /804 System. APCA Journal.
20: 385-390 (1970). * OO

Wilson, W. E., A. Levy, and D. B. Wimmer, A Study of Sulfur
Dioxide in Photochemical Smog II. Effect of Sulfur Dioxide
on Oxidant Formation in Photochemical Smog. APCA Journal.
22: 27-32 (1972).

Miller, D. F., Exploratory Study of Factors Affecting
Aerosol Formation. EPA~650/3-75-002; Contract No. 68-
02-1217 (November, 1974).

Roberts, P. T., and S. K. Friedlander, Photochemical
Aerosol Formation: $02, 1-Heptene and NO, in Ambient
Air. Env. Sci. Tech. (in press).

Cullis, C. F., R. M. Henson, and D. L. Trimm, The Kinetics
of the Homogeneous Gaseous Oxidation of Sulfur Dioxide.

Proc. Roy. Soc. A295: 72-83 (1966).

Vohra, K. G., P. V. N. Nair, and T. S. Muraleedharan,
Possible Role of Singlet Oxygen in an Ion-Induced Reaction
Mechanism of Nucleus Formation by Sulfur Dioxide. Aerosol
Science. 3: 225-236 (1972).

Mulcahy, M. F. R., J. R. Steven, and J. C. Ward, The Kin-
etics of Reaction Between Oxygen Atoms and Sulfur Dioxide:
An Investigation by Electron Spin Resonance Spectrometry.

gd. Phys. Chem. 71: 2124 (1967).

Payne, W. A., L. S. Stiet, and D. D. Davis, A Kinetics
Study of the Reactions of HO, with SO. and NO. J. Amer.
Chem. Soc. 95: 7614 (1973).

Penkett, S. A., and J. A. Garland, Oxidation of Sulfur
Dioxide in Artificial Fogs by Ozone. Tellus. 26: 284-289
(1974).

Armitage, J. W., and C. F. Cullis, Studies of the Reaction
Between Nitrogen Dioxide and Sulfur Dioxide. Combustion
and Flame. 16: 125-130 (1971).

Fair, R. W., and B. A. Thrush, Reactions Between Hydrogen

Atoms and Sulfur Dioxide. Trans. Far. Soc. 65: 1550
(1969).

83


a summary of the limited existing knowledge regarding these pro-
cesses. A third and final report will describe a reactive dis-
persion model developed for use in conjunction with recommended
reaction rate parameters to evaluate the atmospheric impact of
power plant plumes, over moderately large scales of time and

distance.
CLASSIFICATION OF CHEMICAL REACTIONS -- REPORT ORGANIZATION

The reactions by which SO, may be oxidized are generally divided

2
into two groups identified by the terms '‘homogeneous' and 'heter-
ogeneous’'. Homogeneous reactions are those in which all react-
ants (including catalysts) are in the same phase (i.e. gas,
liquid, solid); heterogeneous reactions involve-more than one
phase. In the gas phase, SO, oxidation reactions may be either
homogeneous or, if particulate matter is also present, heterogen-
eous. In the aqueous phase, the significant oxidation reactions
are at least quasi-homogeneous; that is, all the reactants are

in solution, even if the aqueous phase consists of a water sheath
around an aerosol particle. Thus, the important oxidation
reactions to be considered in the atmosphere occur via gas phase
homogeneous and heterogeneous paths and by aqueous phase homogen-
eous paths.

In general terms, the title of a review by Urone and Schroeder ')

essentially summarizes the status of our understanding of So,


FIGURES

Equilibrium Conversion $03 (g) + B20 (g) = #2804 (g)

Dewpoint of H2804 (4)

Sulfur Dioxide Decay Rates in Three Relative
Humidity Ranges (12/19/68, 1/2/69, 1/3/69)

Cross-Plume 03 and SO2 Variations for Various
Downwind Distances

Effect of Acidity on Aqueous Phase Oxidation
Rate

Conversion of SO2 to SO3 in the Presence of
Several Catalysts

Hypothetical Potential Energy - Reaction
Coordinate Diagrams

ii

Page
11

12
19

25

50

68

90


Some of the reactions presented in this discussion are summar-
ized in Table 8. It becomes very apparent when one considers
the role of homogeneous removal paths for SO, that, although
the specific rate data are possibly good only within a factor
of 10, this uncertainty may not be the most serious problem to
assessing the role of homogeneous reaction paths. The various
concentrations of radical species presented in Tables 7 and 8
are not well established for conditions in the troposphere, and
the existing data concerning them may be even less reliable for

conditions in a power plant plume.

One might summarize this section with the conclusion that homo-
geneous SO, oxidation reactions can play a significant role in
atmospheric reactions under certain meteorological conditions,
but one must still question the significance of these reactions
in the power plant plume. At any rate, it is fairly obvious
that we are not confident of just how to apply homogeneous
reaction steps to mechanisms concerning the oxidation of SO,

in plumes.

40


derived from electron spectroscopy for chemical analysis (ESCA)
studies of particulate matter from a premixed hydrocarbon flame
with different SO, levels. The ESCA spectrum of graphite parti-
cles exposed to sO, revealed peaks corresponding to sulfate and
sulfide. Experiments varying the level of SO, added to the
flame as well as the humidity of the air used also showed an
increase in sulfate with increases in each of these parameters.
Although the results of this study show a consistency for carbon
acting as a catalyst, the evidence presented is somewhat limited.
Also, there are no rate data presented in this study from which
one might evaluate the significance of soot in plumes as a
sulfate catalyst.

Corn and Cheng ‘8°? have also carried out studies on the adsorp-
tion of SO, by activated charcoal, as well as by Fe,03, MnO,

and suspended particulate matter. In these experiments, packed
beds of Teflon beads were coated with micron and submicron size
aerosol particles, the beads were packed in a reactor, and pro-
ress of the reaction was followed by continuously monitoring
the effluent S05. Activated charcoal yielded steady-state rates
of conversion, or adsorption, of 0.013 and 0.021 ug SO,/min/mg

charcoal at SO, concentrations of 8.0 and 14.4 ppm, respective-

2
ly. The experiments do not distinguish, however, whether the
50, underwent steady-state conversion in the reactor or whether

there was catalyzed oxidation on the surface.

70


where cy is the concentration and ry the rate of production of
species i in a differential volume element of atmosphere. The
rates tyr of course, will in general be functions of temperature,
relative humidity, insolation, and concentrations of all species.
The solutions which yield values for the c;'s are generally found
by finite difference methods, necessitating rather careful con-
sideration of the technique to be applied, a subject beyond the
scope of this report. The inclusion of chemical reaction mechan-
isms, however, is rather simple since chemical rate equations are
naturally expressed as differential equations, the form required

by the continuity equation.

The program which performs such model calculations can be made
extremely versitile by separating the various functions into dif-
ferent subprograms. For example, in a simple case, the main
program may handle the input/output functions and define such
parameters as the initial calculation point and concentrations,
and step size. A replacable segment may then be called to per-
form the finite difference integration, where the necessary
derivatives are calculated in yet another subprogram. This allows
the examination of different models by the simple replacement of
one subprogram, rather than rewriting the entire code. In such a
configuration, of course, one could also replace the finite dif-
ference method without disturbing the rest of the program, if

such a replacement were desirable.

78


PROPERTY OF

EPA-450/3-76-022 oe
September 1976 METEOROLOGY

SO, OXIDATION IN PLUMES:
A REVIEW AND ASSESSMENT
OF RELEVANT MECHANISTIC

AND RATE STUDIES

U.S. ENVIRONMENTAL PROTECTION AGENCY
Office of Air and Waste Management

Office of Air Quality Planning and Standards
Research Triangle Park, North Carolina 27711



CONTENTS

List of Figures

List of Tables

Sections

I Conclusions

II Introduction

IIL Thermochemistry of SO, Oxidation

Iv Plume Studies

Vv Homogeneous Studies in the Gas Phase
VI Homogeneous Studies in the Aqueous Phase
VII Heterogeneous Studies

VIII Modelling Applications

IX References

x Appendix

13
30

67
77
80

89


Table 5 summarizes some of the SO, oxidation rates and SO, half-

lives reported for measurements made in plumes.

The data are

presented in this table in chronological order of publication.

It is interesting to note in this report the general "slowing

down" of SO, oxidation rates in more recent studies.

The signi-

ficance of this observation is not clear, but it may be indica-

tive of the better, or more reliable, experiments that can be

carried out with today's equipment.

In essence however, truly

reliable estimates for 80, oxidation rates in plumes are still

wanting. None of the data reported to date can be used without

qualification, although this is not meant to fault any of the

studies. Primarily the results show the difficulty of such

TABLE 5. SUMMARY OF SO, OXIDATION

2

RATES IN PLUMES

SO, Oxidation

a/hr tia! hrs R.H., % Reference
Gartrell et al. 0-4 -- 45-70 12(1963)

8-55 -- 73 12 (1963)
Dennis et al. -~ 1.0-2.8 36-53 13 (1969)
Weber -- 0.3-1 -- 16 (1970
Stephens and McCaldin 1.2-2.4 40-80 15 (1971)

0 -- < 35
Coutant et al. -- 1-6 40-90 14(1972)
University of Utah 0.37-0.76 -- 26-50 20 (1975)
Newman et al. 0 -- -- 18 (1975)
Newman et al. -—- 10 -- 17(1975)
Whitby et a7. 1.5-1.8 -- -- 23(1976)

28


58.

59.

60.

él.

62.

63.

64.

65.

66.

67.

68.

69.

70.

McKay, H.A.C., Atmospheric Oxidation of Sulfur Dioxide in
Water Droplets in Presence of Ammonia. Atmos. Env.
5: 7 (1971).

Cheng, R. T., M. Corn, and J. O. Frohliger, Contribution
to the Reaction Kinetics of Water-Soluble Aerosols and
Sulfur Dioxide in Air at ppm Concentrations. Atmos. Env.
5: 987 (1971).

Miller, J. M., and R. G. dePena, Contribution of Scavenged
Sulfur Dioxide to the Sulfate Content of Rain. J. Geophys.
Res. 77: 5905 (1972).

Penkett, S. A., Oxidation of Sulfur, Dioxide and Other
Atmospheric Gases by Ozone in Aqueous Solution. Nature
Phys. Sci. 240: 105 (1972).

Chen, T., and C. H. Barron, Homogeneous Kinetics of Sulfite
Oxidation. Ind. Eng. Chem. Fundam. 11: 466 (1972).

Brimblecombe, P., and D. J. Spedding, Catalytic Oxidation
of Micromolar Aqueous Sulfur Dioxide (I). Oxidation in
Dilute Solutions Containing Iron (III). Atmos. Env.

8: 937 (1974).

Freiberg, J., The Mechanism of Iron Catalyzed Oxidation of
SO in Oxygenated Solutions. Atmos. Env. 9: 661 (1975).

Bigelow, S. L., Catalytic Action on the Speed of Oxidation
of Sodium Sulfite by Oxygen in the Air. 2. Physik. Chem.
26: 493 (1898).

Titoff, A., Contribution to the Understanding of Negative
Catalysis in Homogeneous Systems. 2Z. Physik. Chem.
45: 641 (1903).

Lumiére,A., L. Lumiére, and A. Seyewetz, Changes in
Anhydrous Crystalline Sodium Sulfite by Exposure to Air.
Moniteur Phot. 2: 77 (1908).

Milbauer, J., and J. Pazourek, The Oxidation of Sulfites
in Concentrated Solutions. Chem. Listy. 15: 34 (1921).

Reinders, W., and S. I. Vles, Reaction Velocity of Oxygen
with Solutions of Some Inorganic Salts (III). The
Catalytic Oxidation of Sulfites. Rec. Trav. Chim.

44: 249 (1925).

Backstrom, H.L.J., The Chain Mechanism in the Autoxidation

of Sodium Sulfite Solutions. Z. Physik. Chem. B25: 122
(1934).

85


10.

SECTION IX

REFERENCES

Hales, J. M., Tall Stacks and the Atmospheric Environment.
Battelle-Northwest Laboratories Report to Environmental
Protection Agency, Research Triangle Park, N.C. (January,
1976).

Urone, P., and W. M. Schroeder, SOz in the Atmosphere:
A Wealth of Monitoring Data, But Few Reaction Rate Studies.
Env. Sci. Tech. 3: 436-445 (1969).

Kellogg, W. W., R. D. Cadle, E. R. Allen, A. L. Lazrus,
and G. A. Martell, The Sulfur Cycle. Science. 175:
587-596 (1972).

Cadle, R. D., Formation and Chemical Reaction of Atmospheric
Particles. J. Colloid and Interface Science. 39: 25-41
(1972).

Bufalini, M., Oxidation of Sulfur Dioxide in Polluted
Atmospheres--A Review. Env. Sci. Tech. 5: 685-700
(1971).

Altshuller, A. P., and J. J. Bufalini, Photochemical

Aspects of Air Pollution: A Review. Photochem. Photobiol.
4: 97-146 (1965).

Harrison, H., T. V. Larson, and P. V. Hobbs, Oxidation of
Sulfur Dioxide in the Atmosphere: A Review. Proc. Int'l
Conf. on Environmental Sensing and Assessment, Las Vegas
(September, 1975).

Pierrard, J. M., Environmental Appraisal--Particulate
Matter, Oxides of Sulfur, and Sulfuric Acid. APCA Journal.
19: 632-637 (1969).

JANAF Interim Thermochemical Tables, Thermal Laboratory,
The Dow Chemical Company, Midland, Michigan (1960).

Leppard, W. M., Sulfate Control Technology Assessment,

Phase 1: Literature Search and Analysis. EPA/460/3-75-
002-a; PB-240 995/IWP (November, 1974) 50 p.

80


Thermodynamically, the hydration of SO, to gaseous H,SO0, also

3 4
becomes more favorable as the temperature drops. Figurel,

from Leppard's review (10)

, shows that at temperatures existing
in plumes essentially all the S03 is converted to H,SO, under

equilibrium conditions. This conversion is also achieved rap-
idly: dry sulfur trioxide will form an aerosol almost immedi-

ately upon contact with moist air.

Thermodynamic values for the dewpoint of gaseous H,S0, have been
reviewed by Verhoff and Banchero 1) , Figure 2, again from
Leppard, compares calculated and experimental dewpoint curves.
The rate of condensation is rapid, even at the concentrations

at which SO, or H

3 250

4 is emitted from stacks.
The processes involved in the oxidation of SO, in the aqueous
phase include absorption of SO, and other gases, hydration and
subsequent dissociation of the dissolved species, and oxidation
of sulfite or bisulfite ions. Because these processes are

highly dependent on thermodynamic factors, the appropriate equil-
ibria will be discussed later in conjunction with the relevant
kinetic factors. In general, however, the oxidation process in
the aqueous phase is thermodynamically favored, as it is in the

gas phase.

10


the bisulfite side, decreasing the sulfite concentration. This

effect leads to the expression

a(so, .

-1

where ky is found to be 6.6 + 0.47 wh 2500 This value is

calculated assuming the second dissociation constant of sulfur-

ous acid to be 5 x 107§; more recent research indicates this

-8 (75)

value is actually closer to 6.3 x 10 , leading to a better

“V2 5607},

value of k, = 59M

While the details of the chain mechanism responsible for the

: )
oxidation of S03 are not certain, work by Bassett and Henry ‘76?

and Bassett and parker (50) led to the conclusion that the un-
catalyzed reaction proceeds via ionic complexes such as (0,°

2

SO,)“" and (05°S505) “, The manganese ion-catalyzed reation was

2
3)
postulated to involve an intermediate such as (Oy "Mn ($03) 9) 27
which could rapidly undergo self oxidation and reduction. The
intermediates in the case of cobalt-, nickel-, and iron-catalyzed

oxidation were assumed to be similar but less active complexes.

Quantitative rate measurements were not performed in this study.

One of the first attempts to study the role of SO, oxidation in

air chemistry was done by Junge and Ryan (52) | Their work showed

51


form A and B; that is, the reverse reaction:

c+D7+A+B
is possible. Furthermore, the enthalpy of this reaction is
simply - 4H, and the activation energy E, = E, - Au. (Here the
subscripts f and r denote reference to the forward and reverse

reactions, respectively.)

It should be noted that the discussion above is not limited to
bimolecular reactions. The concepts involved are rather easily
generalized to more complicated systems, and also apply to the
simpler unimolecular reactions. The latter class includes the

extremely important photochemical reactions, which are initiated

by electromagnetic radiation and are involved in many atmospheric

processes.

It must be remembered that the above discussion was a gross simp-
lification of the systems in which we are usually interested, sys-
tems which actually contain large numbers of atoms and molecules.
In particular, atmospheric proceses generally involve reactions of
pollutants at concentrations on the order of parts-per-billion or

10 3 of reactive species

more, meaning some 10 or more molecules/cm
and more than 10? molecules/cm? of inert species. Thus, the

reactant molecules have a distribution of kinetic energies which
are dependent on the temperature, and in such systems, it is pos-

sible to determine a rate at which each species is converted; the rate is

92




ing that assumed quantum yields for SO, removal of 0.01-0.001 are
Much too high, and are actually closer to 1079, On this basis,
Friend would estimate the homogeneous rate of oxidation to be

near zero.

A number of investigators have considered the oxidation of SO,
and subsequent particle formation in the presence of NO.- The

(30)

most pertinent considerations are those of Cox and Jaffee and

Klein (32) | Cox considers the oxidation of SO and of NO, and
their roles in the heteromolecular condensation of H,S0, aerosols.
Using HO as a possible free radical intermediate, the highest

So, conversion rate one might expect from his data is 0.1%/hr.

This estimate is based on the reaction

k
S05 + HO, > S03 + HO ,
where k = 3 x 1072° om3 molecule"! secTM! and the HO, and SO, con-

2 2
centrations are 5 x 108 molecules/cm? and 30 ppb respectively.

Cox's analysis shows that the rate of formation of H,S0, in urban
air is sufficient for hetermolecular nucleation to form aqueous
sulfuric acid aerosols, while in background areas SO, oxidation

products are removed mainly by condensation on existing aerosol.

°
Jaffee and Klein irradiated NO, in the presence of SO, at 3660 A,a

wavelength at which ce) is not excited, but NO, is photolyzed to NO

k *
and 0. The rate constant for the reaction SO, + ie) ot SO, is

reported to be 1.1 x 107 2 mole! sec. When combined with the

34


TABLE 10. (CONTINUED)

RATE COEFFICIENTS:

AUTHOR TYPE OF MECHANISM ELEMENTS OF MECHANISM AND/OR_ EXPRESSIONS COMMENTS
~ - (05) ‘
Penkett S02 oxidation by ozone | HSO3” + 03 + B50,” + 07 3) ~) | 9.62¢: S02 oxidation
1972 ° ‘ “ae = ¥2(03) (S037) | rate extrapolated from
data; much faster than
(61) ky = 3.32 + .13 x Scott and Hobbs
105ymts~1
202 conversson rate
% f228/main
: ; e105)
Chen & Barron | sulfite oxidation free radical mechanism; 2 Mt)
1972 catalyzed by cobalt Co(t2I) reduced gould not determine |
Sons xe[co(p0)63+]/2 | SPecstae value for
(62.
) [so3=]3/2
Penkett and HSO3” + 03 + HSO4” + 02 81509)

Garland
1974

(42)

$0 oxidation by ozone

~ art + ¥[5004]
k= 4.18 x 1074 +
+ :1.77 [Ht)1/2 secm}

pH range from 4 to 7
10°C; «1 ppm S02,
+05 ppm 03 in fog water

Brimblecompe & $0) oxidation by 0; complex - aisttvy) . Possibility of Fe(III)
Spedding with trace Fe catalyst SUS EOL wineczazy) | [oncaminacion discussed
1574 {s(zv))
(63) k= 100 wel sol
802 conversion rate
‘v3.2%/day in fog
assuming 28 ug/m? soz
and 107 M Fe(III}
Freiberg | $02 oxidation catalyzed | complex 6 4804*] Kg2lHp$03]2| Rate ancreases rapidly
19 by’Fe ie KoKs*'H2S031") with RH and decreases
by about one order of
(64) {Fe3+}/{H7]2 | magnatude wath S$°C
increase in temperature.
ke = 18t gissocaation

‘constant of #2503

45a


Time in seconds.

FIGURE 5, EFFECT OF ACIDITY @9)

Curves 1 to 8 refer respectively to 0.02 molar solutions of sodium
sulfite to which 0, 0.0032, 0.0065, 0.0097, 0.0130, 0.0162, 0.0195

and 0.0325 mole per liter of hydrogen ion have been added as
sulfuric acid.

>

50


been either ammonia or a metal salt or oxide, most likely Fe(III)
or Mn(II). (Some investigations have postulated reaction in the
aqueous phase of a deliquescent aerosol particle; in mechanisms of
this type, the classifications homogeneous and heterogeneous lose
much of their meaning, and these reactions are rather arbitrarily
included in this section of the report.) Aqueous phase oxidation
of SO, by 03 has recently been measured and reported. A very lim-
ited amount of work has been done on photochemically induced oxi-
dation in the aqueous phase, put it is generally discussed as a
side reaction in (and dependent upon) the photoxidation of olefins
and reactive hydrocarbons. Because of the low output of these
species from power plants, it is expected that photochemical oxi-
dation of SO, in the aqueous phase is not a significant source of
sulfates in power plant plumes.

The following discussion of the oxidation of SO, by oO, in the

2
aqueous phase will assume (unless otherwise stated) that the

reaction is either

HSO 4

3 1/2 O, > HSO

or

S03 + 1/2 2 > SO

4 .
It will be clear from the context which reaction is being discus-
sed. This section of the report will consider first some general
work on SO. oxidation in the aqueous phase, and then the specific

effects of ammonia and metal salts. Finally, oxidation by ozone

and photochemical processes will be discussed briefly.

46


n (92)

Smith, Wagman and Fis. report sorption studies of So, on

Fe,0,, Al,0, and platinum, using particles ranging in size from

about 0.01 to 0.1 micron diameter. Using a °>

80, technique, they
developed sorption isotherms. Their technique may prove useful in
Gistinguishing layers of preferential chemical adsorption from
multilayers of physical adsorption. They observe significant

adsorption on all three surfaces, but do not measure oxidation

rates in these studies.

Some basic studies on the adsorption and conversion of SO, on
CaO and MgO have been carried out as part of efforts to remove
SO, from power plant stack gases by adding limestone or dolomite.

Two studies from investigators at New York University (92/93

have
been concerned with infrared examination of SO, adsorbed on CaO
and on MgO. Sulfites are produced on the surface and subsequent-

ly oxidize to sulfates on heating in air.

Lunsford has also studied the SO,-MgO system, both at high temp-
eratures and at ambient temperatures ‘94 , The low temperature study,
although it provides no rate data, has more potential relevance

to plume chemistry. Samples of MgO and Mg (OH) 5 were exposed to

SO, at 25 torr, and electron paramagnetic resonance (EPR) and
infrared (IR) spectra were obtained. The results showed that
heterogeneous oxidation occurred on the MgO particles. Sulfite
ions and monodentate sulfite complexes were identified. The

investigators suggest that water adsorbed with the SO, increases

74


45.

46.

47.

48.

49.

50.

51.

52.

53.

54.

55.

56.

57.

Falk, M., and P. A. Giguere, On the Nature of Sulfurous
Acid. Can. J. Chem. 36: 1121-1125 (1958).

Wang, J. C., and D. M. Himmelblau, A Kinetic Study of
Sulfur Dioxide in Aqueous Solution with Radioactive
Tracers. Am. Inst. Chem. Eng. J. 10: 574-580 (1964).

Eigen, M., K. Kustin, and G. Maass, The Speed of Hydration
of SOj in Water Solutions. Z. Physik. Chem. Neue Folge.
30: 130 (1961).

Yue, G. K., C. S. Kiang, V. A. Mohnen, and E. Danielsen,
The Interaction of Atmospheric Sulfur Compounds with Cloud
and Precipitation Elements. State University of New York
(Albany) Interim Report to National Science Foundation,
143 p. (May, 1975).

Fuller, E. C., and R. H. Crist, The Rate of Oxidation of
Sulfite Ions by Oxygen. J. Am. Chem. Soc. 63: 644 (1941).

Bassett, H., and W. G. Parker, Oxidation of Sulfurous
Acid, J. Chem. Soc. 1951:1540.

Junge, C. E., and T. G. Ryan, Study of Sulfur Dioxide
Oxidation in Solution and Its Role in Atmospheric Chem-
istry. Quart. J. Roy. Meteor. Soc. 84: 46 (1958).

van den Heuvel, A. P., and B. J. Mason, The Formation of
Ammonium Sulfate in Water Droplets Exposed to Gaseous
Sulfur Dioxide and Ammonia, Quart. J. Roy. Meteor. Soc.
89: 271 (1963).

Espenson, J. H., and H. Taube, Tracer Experiments with
Ozone as Oxidizing Agent in Aqueous Solution. Inorg.
Chem. 4: 704 (1965).

Scott, W. D., and P. V. Hobbs, The Formation of Sulfate
in Water Droplets. J. Atmos. Sci. 24: 54 (1967).

Beilke, S., and H. W. Georgii, Incorporation of Sulfur
Dioxide into Fog and Rain Droplets. Tellus. 20: 435-442
(1968).

Foster, P. M., Oxidation of Sulfur Dioxide in Power Station
Plumes. Atmos. Env. 3: 157 (1969).

Matteson, M. J., W. Stéber, and H. Luther, Kinetics of

the Oxidation of Sulfur Dioxide by Aerosols of Manganese
Sulfate. Ind. Eng. Chem. 4: 677 (1969).

84


*IM/ZT00"O > 89781 Oxy

“ptt OF X S°Z =H ‘49p20 puvdos su passordKe UMSUHD Geynoojowsny,
” at? oT 389 (H) ° 4 yeOl EE cat) Zosn = (HH) n+ 20g
ITZ ¥ OOE 09 ¥ OOL WoXY parefodearxe ¥ ° 1c-O8 * 67%
1% z ry mecbl XZ > ‘ 2
xUTe xT =
Te cot? 200 * oT = (onl ry pol * 1-9 on + fos = “on + os
a y= ud‘ “ Sz°0 OLX az
zy 9a nd wo ort oz = tol “Sz Sirol x BZ otn +) 20 + “og = (on 4) £0 + 20s
Tz . 910 ot * 6
87 . z soo LOE xe ¢ a2
W a> yor x s = (on) ot OLX 6 on + fos = Son + 208
og st
iv . ara zy-08 * Le
L : 6°8 aot ¥ Sz
1z 0 6°e eyo x11 ‘ 2
9% cm ,0r = [out] oT co es-8°0) ait) fosn = (at) on + “os
" 0 01x 6't
. 0 ; yor x ZT
“ 40 y cro! * ze
ut cm (oT *S = [Cao] 100° zy0l * 8°1 € z
07 wo For x s = [(a,)o] 100°0 Tol * 61 at) “os = (nt) 0 + 20s
cs Ic Pea
6t = ((80,)701 To" gi-Ol 0 + fos = (8vz) Zo + Zos
tt T00"0 gz-01 * 6°S
Be s0°0 g7-0l * 6% fos = Zo z/T + 20s
SDT TST GUSTIS FISTS AIST OM TH ~ - -
* (aqewyxoadde) jueqsuog aby
_ aru
M00 FL

‘Twaowau 70S 40 SHLWE GULVWILSA 8 TTOVL

41


This report is issued by the Environmental Protection Agency to report
technical data of interest to a limited number of readers. Copies are
available free of charge to Federal employees, current contractors and
grantees, and nonprofit organizations - in limited quantities - from the
Library Services Office (MD35), Research Triangle Park, North Carolina
27711; or, for a fee, from the National Technical Information Service,
5285 Port Royal Road, Springfield, Virginia 22161.

This report was furnished to the Environmental Protection Agency by
Battelle Pacific Northwest Laboratories, Richland, Washington 99352,

in fulfillment of Contract No. 68-02-1982, Program Element No. 2AC129.
The contents of this report are reproduced herein as received from
Battelle Pacific Northwest Laboratories. The opinions, findings, and
conclusions expressed are those of the author and not necessarily those
of the Environmental Protection Agency. Mention of company or product
names is not to be considered as an endorsement by the Environmental
Protection Agency.

Publication No. EPA-450/3-76-022

ii


mechanism similar to that above is responsible, with the initi-

ator being FeOOH:

3
and the Fe(III) may be regenerated by oxidation of Fe’t,

= +
$0," + FeOOH + 3H” + Fe 2* + 2H,0 + -SO

- 2+ 3+

+ Fe + 0, + Fe

allowing oxidation of large amounts of SO, by very small amounts

of Fe(III). (The hydrozylated Fe(III) species FeOOH has a

3+

higher oxidation-reduction potential than Fe (0.908 V com-

pared with 0.77 yy (75480) | and pH values in the range pH 4
to pH 5 tend to favor formation of the hydroxylated species.

Thus, FeOOH appears to be a better oxidizing agent than Fe?*

for the production of “S03 radicals).

In another study of the catalytic effect of iron, Freiberg (64/81)

obtains a rate expression:

,

d1S0q ] _ KpK(Fe?*] [H,80,17

dt tnt73

where K is the first dissociation constant for H,S0, and Kp is
a rather complicated function of [05], Fe 24), and several

equilibrium constants; the expression fits the published data of

63


TABLE 11. EFFECT OF ACIDITY ON REACTION veLocrry ‘49)

el

Hydrogen Total Sulfite
ion Sulfite ion
added, Time, Conen., Conen., . «
m/l sec. n/l ~ m/l pH 1
0.0032 ie) 0.0164 0.0132 5.92 1.9%
20 -0129 -0097 5.77 5.4
40 -0096 0064 5.60 5.9
60 +0072 -0040 5.40 6.3
80 +0055 -0023 5.15
«0065 0 -0160 -0095 5.47 6.3
20 0121 +0056 5.24 7.0
40 +0094 .0029 4.96 6.6
60 .0078 +0013 4.60 6.9
80 0069 ~0004 4.09
-0097 0 -0145 -0048 5.00 6.8
20 20119 +0022 4.66 7.2
40 +0104 -0007 4.16
+0130 0 +0160 .0030 4.66 7.6
20 +0139 -0009 4.14 6.9
40 -0131 0001 3.19
-0162 0 +0174 0012 4.17 6.1
20 +0164 -0002 3.40

Av = 6.6 = 0.47

(a) This figure was not used in calculating ky (average).

49


TECHNICAL REPORT DATA

(Please read Instructions on the reverse before completing)

7. REPORT NO, 2
EPA-450/3-76-022

3, RECIPIENT'S ACCESSIONNO.

fs. TITLE AND SUBTITLE
SO2 Oxidation in Plumes: A Review and Assessment
of Relevant Mechanistic and Rate Studies

Is. REPORT DATE Date of Approval
September 1976

6. PERFORMING ORGANIZATION CODE

7. AUTHOR(S) .
° A. Levy (Battelle, Columbus Laboratories),

D. R. Drewes, and J. M. Hales

8. PERFORMING ORGANIZATION REPORT NO.

9. PERFORMING ORGANIZATION NAME AND ADDRESS:
Battelle, Pacific Northwest Laboratories
Battelle Boulevard
Richland, Washington 99352

10. PROGRAM ELEMENT NO.
2ac 129

77. CONTRACT/GRANT NO.
68-02-1982

12. SPONSORING AGENCY NAME AND ADORESS
Office of Air Quality Planning and Standards
Environmental Protection Agency
Research Triangle Park, North Carolina 27711

13. TYPE OF REPORT AND PERIOD COVERED
Final

14, SPONSORING AGENCY CODE

15, SUPPLEMENTARY NOTES.

76, ABSTRACT

is reviewed.

oxidation of SO.

significant processes.

by maintaining a high pH.

The scientific literature pertaining to the oxidation of S02 in power plant plumes
Aqueous phase, homogeneous gas phase, and heterogeneous gas phase
mechanisms are considered, as are actual plume studies.
over a wide range, which is not totally unexpected due to the highly complex
nature of the oxidation process, and some general conclusions can be drawn:

The reported rates vary

1) Recent plume studies, in general, indicate lower rates than earlier ones, and
also suggest a coincidence between the reappearance of ozone in the plume and

2) Gas phase studies indicate homogeneous reaction of S02 with OH radicals and
heterogeneous reactions catalyzed by lead and iron to be perhaps the most

3) In the aqueous phase, the reaction is most effectively catalyzed by iron
and manganese, and ammonia plays an important role in promoting the oxidation

H7. KEY WORDS AND DOCUMENT ANALYSIS

2. DESCRIPTORS

Ib. IDENTIFIERS/OPEN ENDED TERMS |c.

COSATI Field/Group

Electric Power Plants
Reaction Kinetics
Sulfur Dioxide
Mathematical Models

13. DISTRIBUTION STATEMENT.

19. SECURITY CLASS (This Report)

21. NO. OF PAGES

Unclassified 100
Release Unlimited 20. SECURITY CLASS (This page) 22, PRICE
Unclassified

EPA Form 2220-1 (9-73)



oxidation rates, but because it introduces rather different
thoughts on the interaction of SO, in plumes. They col-
lected extensive data on NO, NOo, and 03 and less extensive data
on SO in the plume of a 1000-MW power plant firing a mixture of
oil and coal. The significant observation in this study was the
“ozone bulge" depicted in Figure 4. In these profiles the plume
boundaries are well characterized by the drop in the ozone
concentration. This is readily accounted for by the rapid

reaction between NO and 03 (NO + 03> NO, + 05). At 50 km down-

2
wind, however, one observes an increase in ozone (the "bulge")
coincident with low levels of S05 measured as "plume so,"

The authors suggest a series of rather questionable reaction
steps to explain the bulge. Of interest to this review, how-
ever, is the fact that these reaction steps, if correct, would

probably lead to H,SO and aerosol products which would have

4
some bearing on the rate of oxidation in the atmosphere. The

general reaction sequence suggested in this work is:

°

2
SO + OH + M ——> HSO,———> HSO,
NO, + HSO, < NO
ESO.
2NO
2N0. + HSO, <

L,, 3

24


leaving the stack of a power plant in Frankfurt am Main, he then
observed CO/S05 ratios at various distances from the plant. His
results only allow for generalized statements on rates of oxida-
tion. He reports rather short half-lives (20 minutes to 1 hour)
depending on meteorological conditions (which are not well-
defined). He also reports an increase in 4C0,/480, with increas-
ing relative humidity, indicative of an increasing oxidation rate
(here, A indicates the difference between the peak and background

levels).

The Brookhaven National Laboratory (BNL) has been carrying out
fairly extensive studies of So, oxidation in plumes of oi (17)

and coal-fireg ‘18) power plants. The BNL studies were carried

out with the aid of single engine aircraft outfitted with a high
volume sampler for collecting and measuring sulfate and with S05
absorption scrubbers for applying the BNL isotopic ratio tech-
nigue (19) | The particulate sulfate catches were too small to be
adequate for accurate analysis. The isotopic ratio technique
depends on measuring the changes in the ratio of the two most
abundant isotopes of sulfur ,?*s/?4s. Since the technique measures
the sulfur originating with the fuel, it can theoretically

discriminate between sources, and thus lends itself to the study

of a specific source.

20


TABLE 6, ESTIMATED RATES OF THEORETICALLY POSSIBLE
HOMOGENEOUS REMOVAL PATHS FOR $Q2 IN A
SIMULATED POLLUTED ATMOSPHERE @©

a

Approximate
ABo gg, Rate, %
kcal/mole per hr
50, + 1/2 9, + Sunlight - 50, -24 < 0.021
oC) + $0, +M~ $0, +M -83 0.014
03 + 80, ~ 80, + o, 56 0.00
No, + So, = SO, + NO -10 0.00
NO, + 50, = SO, + No, -33 0.00
NO. + SO, - 80, + N,0, -24 0.00
ie}
7 3N
CH, —— CH, + So, “ 80, + 2CH,0 -81 < 0.4-3.0
CH,00 + SO, - 50, + cH, ~ -117 < 0.4-3.0
CH, = 040 + SO, - So, + CHAO ~ -85
HO, + 80, - HO + SO, -19 0.85
= ; - ?
HO S05, < +25
CHAO, + So, _ CH,0 + 50, -30 ~ 0.16
- : - 2
CH,0,S0, < +25
HO + So, - HOSO,, ~ -82 ~ 0.23
CH,0 + SO, - CH,0SO,, ~ -73 ~0.48

Total potential rate of conversion of So, to SO, (or Sulfates)

= 1.7-4.7% per hour.

3

32


10 ug NE,/m°. Keeping in mind that the data are only semi-
quantitative, it is still apparent that the process is of impor-
tance in the oxidation of S04 in droplets as well as in bulk
solution.

A much-quoted theoretical investigation by Scott and Hobbs (54)
assumes equilibrium is maintained between gaseous and dissolved
S05, NH3, and COnr and between the ions produced in the dissoci-
ation of the dissolved species:

2°H20 K

yi

> .
809 (g) + H0 + so hs = [so H,01/P oo

2

S0,*H,0 + HsO,” +H” Ki, = [HS0, 1 [#"]/[S0,°H,0]
Hso, = = so, + Ky, = [80°] [#7]/[#s0,7]
NH3(g) + #20 2 NH,-HZ0 "Kya = [NH5*H201/Py
NH,"H,O 7 NH,” + onT Ky, = (NH,7] [0H ]/(NH,-#,0]
CO (g) + Hy0 2 c0,*H,0 Knc = [C0,*H,01/Peq
CO5*H,0 z HCO,” + Ht Kiet [HCO37] (HI /{C0,°#,01
Heo, = co,” +H” - Rog = (co, ] [#7] / HCO, 7]
H,o 7 Ht + on K, = [H*) [0H]

Using the appropriate equilibrium expressions and constants, it
is possible to derive an expression for the sulfite concentra-

tion as a function only of (Ht), which can in turn be determined

55


Active sites on the surface of the particles become occupied by
the products of the reaction, and thus are not available for

further reaction. The capacity of the Fe,0, particles to oxid-
ize 80, in air was found to be 62.6 ug/mg Fe,03. The rate con-

stant for the heterogeneous reaction was 9.4 x 1073 ppm min7?

(6.3 x 10728 om? mote?

sec"), SO concentrations in this study
were varied from 4.7-18.8 ppm and relative humidity was varied
from 50-94 percent. The rate constant did not appear to be cor-

related with so, or relative humidity.

Urone et a1{99) examined the reactions of S05 in the presence of
hydrocarbons, nitrogen dioxide, moisture, particulate matter and
ultraviolet radiation. Gaseous mixtures did not react in the
dark and SO, oxidation was of the order of 0.1%/hr under ultra-
violet irradiation. In the presence of particulate matter, the

reaction was significant in some instances.

Table 14 presents a summary compilation of SO conversion rates
drawn from the Urone study. SO concentrations in this study
were varied from 8 - 14 ppm; particulate loadings were 16 -
30 mg. The results show high reactivity with the iron, lead and
calcium oxides and low reactivity with Nacl, Caco3, A1,0, and
V305- The results compare well qualitatively with those of

(88)

Corn and Cheng 1 who reported high reactivity for Fe,03 and

low reactivity for caco, and V,0,.

72


These experiments are reported here with no further comment.
Some other work by this group is discussed in the section on
homogeneous gas-phase reactions.

Whitby et at. ‘23)

conducted sampling flights through the plume of
a large power plant near St. Louis, Missouri, as part of the Envi-
ronmental Protection Agency's Project MISTT (Midwest Interstate
Sulfur Transport and Transformation). Using the data from one
particularly good run in 1974, they estimate SO conversion rates
of 1.5%/hr between 10 and 21 km, and 1.8%/hr between 21 and 32 km

downwind. The relative humidity was 75% and the sulfur flow

appeared to be constant at 4.08 kg/sec.

Roberts and Friedlander (24)

have recently carried out some calcu-
lations which permit estimates of gas-to-particle conversion rates
for sulfur in trajectories in the Los Angeles basin. These analy-
ses are not based specifically on power plant plume trajectories,
but do provide interesting results of pertinence to the oxidation
of sulfur dioxide. The authors used total filter and cascade
impactor aerosol samples from several sites in the Los Angeles
basin. Depending on source and wind direction, S0,/Total S ratios
ranged from 0.88 to 0.98. Estimates of pseudo-first-order rate
constants, which appear to be rather high, are presented in Table
4. The accuracy of these data, like all data of this type, is
strongly dependent on the method of calculation, the measurement
technique, and such complex mechanistic parameters as ozone, free

radicals, and humidity.

26


Exothermic Reaction

=a _
Es
A+B
E
AH r
aw C+D

Reaction Coordinate

Endothermic Reaction

E,
ce, [+ cto
ate | J 4H
Ww

t

Reaction Coordinate

FIGURE 7. HYPOTHETICAL POTENTIAL ENERGY--
REACTION COORDINATE DIAGRAMS

90


oxidation chemistry: "SO, in the Atmosphere: A Wealth of Mon-
itoring Data but Few Reaction Rate Studies". It is of interest
that despite numerous studies on both the heterogeneous and
homogeneous aspects of SO, oxidation, there is no consensus on
the predominant path of this reaction in the atmosphere in
general, or in plumes in particular. There are more studies
supporting the predominance of heterogeneous paths, but some
recent studies seem to suggest homogeneous rates greater than
two percent per hour, higher than previously observed. In
addition, the difficulties inherent in assessing the importance
of the various paths, especially the aqueous phase reactions,
are compounded by an acute dependence on meteorological factors.
In this report, the oxidation of SO, is reviewed as follows:

Thermochemistry

Plume studies

Homogeneous studies in the gas phase

Homogeneous studies in the aqueous phase

Heterogeneous studies ,

Modelling applications.

The Appendix contains a brief review of thermochemical kinetics
and definitions of the basic chemical terms which will be used
throughout the report. It is included for the sake of complete-

ness and also as an aid for the non-chemist.

The literature concerning various aspects of the oxidation of

S05 has been reviewed by several authors (Urone and Schroeder

7


TABLE 14. SUMMARY OF SO, REACTIONS ON
VARIOUS PARTICULATE SPECIES (90)

Particulate Species % SO, Reacted/min
Caco, -004-.07
Crj03 -009
V205 0-.015
Nacl 0-.02
Al,0,3 -04

cao 1.8-3.0
A1,0,/Ca0 2.4-2.6
Pbo 1.9

PbO 5.9

Fe30, 4.2

Fe,03 4.5

73


The actual inclusion of chemical rate data into a plume model,
then, is not a great problem. The major obstacle to the use of
such models is the difficulty of finding reliable data to use
for input, since the result of any calculation can be no better
than the assumptions and data upon which it was based. As the
preceding sections point out, data available concerning SO,
oxidation rates are not extensive, especially for heterogeneous
gas phase processes, a fact which severly limits the immediate
use of models as predictive tools. There are perhaps more impor-
tant uses for them as diagnostic tools in the intercomparison of
various mechanisms. It can be hoped that future studies will
improve our understanding of the important S50, oxidation process
and allow more informative diagnostic plume modelling than pre-

sently possible.

79


tion rate increasing downwind as ozone returns to ambi-
ent levels after having been depleted by reaction with
NO in the early plume.

5) Hydrocarbons may promote SO, oxidation via photochemical

smog processes.

Homogeneous Gas Phase Studies

1) Rates are expected to be from 1.6 to 13.3%/hr.
2) Evidence indicates that the reaction of S05 with OH may
be the most significant process in the atmosphere, but

OH levels in plumes are not known precisely.

Aqueous Phase Studies

1) Reported rates range from 0 to 20%/hr.

2) Dissolved manganese and iron are the most efficient
catalysts, and both species are expected to be present
in plumes.

3) Hydration of aerosols, particularly manganese sulfate,
can form an aqueous phase capable of oxidizing SO, at

an appreciable rate.

Heterogeneous Studies

1) Oxidation rates range from 0 to 6%/hr.
2) Lead and iron are efficient catalysts promoting high
reaction rates; vanadium, surprisingly, appears to be an

inefficient catalyst in plumes.


SECTION VIII
MODELLING APPLICATIONS

As suggested in the previous sections of this report, the measure-
ment of SO, removal and/or sulfate formation rates, whether in the
field or in the laboratory, presents a variety of problems. While
the details of using these data in a model to predict sulfate for-
Mation are likewise formidable, the technique itself is at least
conceptually uncomplicated. The third report to be prepared for
this project will be the more detailed documentation of an actual
model, STRAC (Source-Transport-Receptor Analysis Code), developed

for this purpose.

The technique, in general, must create the model-world analog of
a real-world situation through a suitable mathematical formalism.
The practical limitations dictated by computer capacity and cost,
as well as the more philosophical ones imposed by an imperfect
knowledge of nature, of course, determine how good the analogy is.
A model of a reactive power plant plume must include a means for
calculating the spatial distributions of pollutants, and it is
this portion of the modelling which involves chemical reaction
data. Specifically, the model is defined by the general continu-

ity equation for each of n chemical species (S03, S03, NO, 03-+-)
= .
— = ~¥s(vie5) + ry (i=1,2,...n) ,

77


94..

95.

96.

97.

98.

Lin, M. J., and J. H. Lunsford, Photooxidation of Sulfur
Dioxide on the Surface of Magnesium Oxide. J. Phys. Chem.
79: 892 (1975).

Li, K., R. R. Rothfus, and A. H. Adeny, Effect of Macro-
scopic Properties of Manganese Oxides on Absorption of
Sulfur Dioxide. Env. Sci. Tech. 2: 619-621 (1968).

Happel, J., and M. A. Hnatow, Catalytic Oxidation of
Sulfur Dioxide Using Isotopic Tracers. PB No. 224-305,
(EPA-650/2-73-020) 33 p. (August, 1973).

Opferkuch, R. E., et al, Applicability of Catalytic
Oxidation to the Development of New Processes for Removing
SO2 from Flue Gases. Vol. III. (PB 198-810) U. S. Dept.
of Health, Edication, and Welfare, Cincinnati, Ohio

260 p. (August, 1970).

Merryman, E. L., and A. Levy, Sulfur Trioxide Flame
Chemistry. XIII Symposium (Int'l) on Combustion, Univ.
of Utah, The Combustion Institute, Pittsburgh, Pa.

p. 427 (August 28, 1970).

88


the work indicated that the reaction should proceed faster as
the temperature is lowered, a feature which had been noted by
other investigators and may be attributed to increased gas
solubility at lower temperatures. Some of McKay's data suggest

rates of oxidation of about 13 percent per hour.

There is still no concensus about the "best" value for the ox-
idation rate for this system. Data taken by Miller and de

Pena (60) fit a curve corresponding to k ~ 0.003 sect, very
close to the value used by Scott and Hobbs. Some of the Miller and
de Pena curves, however, suggest oxidation of only 0.1 percent

per hour.

The relevance of the entire SO,-NH,-water system to the case of
power plant plumes is quite uncertain. Reported SO, oxidation
rates attributable to ammonia catalysis vary widely, from 0.1 -
13 percent per hour. The NEG concentrations and pi calculated for
the system, regardless of k, are much higher than measurements
indicate. One possible explanation for this involves the sol-
ubility of NH in water: if it is actually a good deal less
than the models assume it is, both [NEG] and pH predictions
could be lowered. This would mean lower 1so,] and thus, for
the Scott and Hobbs-type analysis, the rate would be lower. For
an analysis using McKay's formula for k, however, the lower pH
also means a larger value for k, and for reasonable ambient

levels of SO, and NH3, the rate will actually increase.

57


TABLE 12. THE EFFECT.OF VARIOUS SOLUTIONS ON SULFATE
FORMATION (51)

Solution of so,” Concentration
51 x 104 g/m SO, and after 3 hr
MnCl, (1 yg/em?) 329 ug/em®
CuCl, " 199

FeCl, 167

Coll, " 49

NH,OH 49

NaCl " 4

dist. HAO 3

53


TABLE 3. SO, OXIDATION STUDIES COLBERT STEAM PLANT prog 1?)

rr

Travel from Point of Emission Relative
‘ : . Humidity S02
Date Sample Time Distance in Plume Oxidation
1960 No. (min) (miles) (Percent) (Percent)
“Low Rates"

8/2 1 5 225-1 0

2 5 125-1 0

3 5 225-1 1.20

4 15 1-1.5 )

9/2 1 30 2-3 3.70

2 78 8 2.20

10/14 1 12 251.5 62 2.15
2 60 5-6 54 3.23

10/26 1 6 +25-1.25 45 1.50
2 84 8-9 48 2.70

10/28 1 12 25-1.5 68 1.10
2 84 8-9 70 4.10

“High Rates"

5/3 1 13 1.1 13.80
2 13 lL 10.00
3 13 1.1 19.20
8/19 1 108 8-10 55.50
2 23 275-2 8.00
10/11 1 12 ~5-1.5 74 21.60
2 96 8 73 32.00

16


TABLE 10. SUMMARY OF AQUEOUS PHASE S02 OXIDATION MECHANISMS

AUTHOR, _ I

TYPE OF MECHANISM

ELEMENTS OF MECHANISM

RATE COEFFICIENTS
AND/OR EXPRESSIONS

COMMENTS

Fuller ané crast |
1941

(49)

Sulfite oxidation by 02
with no catalyst;

cu2* catalyst; mannitol
inhibitor

very long chain reaction

a(0,) -

omge = & (50377)

k= .013 + 2,5(Cu2*)
with cu2* catalyst;

k= 1013 + 6.6 (Ht)2/?
with acid added

25°C; rate coefficient
wath acid added assumes
second dissociation of
H2S03 equals 5 x 10-6
(too large)

Bassett and
Parker
1941

(50)

Sulfurous acid oxida-
tion by metal salts
and 02

uncatalyzed: formation of

such as [02°S03]~7;
formation of com-
such as [02-Mn(S03)2]~

plexe
and rapid oxidation

Junge and Ryan
1958

(51)

$02 oxidation catalyzed
by Feé*, wath and with-
out ammonia

conversion rate
= 1.8 x 1074 8/man

sulfate formation
asymptotic to a certain
pH; maximum formation
varies linearly with
SO) partial pressure

Van den Heuvel

$02 oxidation catalyzed

2802 + 02 + 2803

conversion rate

extrapolation to atmo-

and Mason by “NH . . = | = 2.58/mn spheric conditions use¢
13e2 3 S03 + H20 + H2S0q ~ 2H*+S04) to estimate conversion
(52) NH3 + #20 + NH,OH Z NH,*+OH" rate

Espensen and S02 oxidation by Ozone | S02 + 03 + HzO > HSOg” + not determined

‘Taube Heo.
1965 2

(53)

(in acid solution)
$03" + 03 + S04" + 02
(in basic solution)

Scott and Hobbs
1967

S02 oxadation catalyzed
by“ NH3

S02, NH3,
with ions

C02 in equilibrium
and gas

4] 50,"
alee] x(s03"]

k = .0017 secTM)

25°C only: k estimated
from Van den Heuvel
and Mason

484) $03" + 1/2 02 + SO4TM
Beilke and Washout and rainout of assumed anstantaneous | calculated S04" concen-
Georgia $02 and sulfate aerosols) oxidation of all S02 trations from 0.3 to
1968 incorporated into 36.5 mg/£ depending on
($8) droplet rain rate and assumed
80) gas phase concentra-
tion
Foster, S02 oxidation catalyzed | 2502 + 220 + 02 S02 conversion rate = | theoretical study; rat
6 by*metal salts for Mn and Fe are func-
(56) gatelyst, 25504 1pgtimin for tne tions of many factors:
:15 to 1.5t/min for Fe
for typical input rate far Fe catalyzed
paranctese oxidation is pH depend-
ent
: (505)
Matteson, S07 oxadation catalyzed | Mn2* + so, 2 Mn-s022+ ~ at = y[un?*]3 | nerigibie soy”
Stober, and by’metal salts formation for RH <958
Luther 2Mn-$022+ + 02 ky = 2.4 x 105 mri s-) | similar mechanism may be
1969 . 1 responsible for catalysis
(57) = [(re-s02?*) 2°02 by other metal salts
’
= 2mn-s032+
Mn-S0324 + 420 7 Mn?* + HSO4”
+e
+ Bt = 250,
4/804 “1
Kay $02 oxidation catalyzed | same as Scott and Hobbs pit = k[s03") considered variation
1972 by*wHy with temperature: k esti-
w= (.013 + 5918+) 2) mated from Fuller and
(58) seert

= .032 secwl at
neutrality (25°C)

Crist; found large nega-
tive temperature rate
correlation

Cheng, Corn, &
Frohlinger

(59)

S02 oxadation catalyzed
by metal salts

2802 + 220 + 07 Sabaivse.

28250,

802 conversion rate
~.O30/min wath Mn2*
levels typical of urban
industrial atmospher;
~.338/min with levels
typical of plume from
coal powered plant

oxidation rate estimated
by extrapolation to
atmospheric conditions

Miller and

DeFena
1972
(60)

$02 oxidation catalyzed
by NH3

Similar to Scott and Hobbs

4'so,"
it = k|s03"]

k x .003 sec7) (25°C)

k estimated by fit of
experimental data to
theoretical curves

45


The earliest definitive work on the process was reported by
Fuller and crist (49) in 1941. In the absence of added catalyst,
their data showed the reaction $0,” + 1/2 05 Pa SO,” to be first
order with rate constant k, = 0.013 + 0.0015 sec}, the work was
performed at 25°C with pure oxygen atmospheres. When the inhibi-

tor mannitol was added, the data were found to fit an expression

similar to B&ackstrém's, and over a wider range of inhibitor con-

centrations:
- —A =
- = gam Kz (593)
where A is about 1075, The catalytic effect of cut was also

investigated, and it was found that the data fit an expression

of the form

a(S0, } ++ =
_ 3 = [s + k3 (Cu ) (SO, )
dt
and k3 was determined to be 2.5 + 0.33 x 10% wtsect. The ex-
treme sensitivity of the reaction to catalysis by cutt can be

seen by noting that at concentrations greater 5 x 1079 M,

k3(Cu"*) > k}- The effect on pH on the oxidation was also invest-
igated, and as noted in previous (and subsequent) work, the
oxidation rate increased with increasing pH (see Table 11 and
Figure 5). This is consistent with the assumption that the sul-
fite rather than the bisulfite is the species being oxidized;

->
=

the addition of HT shifts the equilibrium (HSO3 HY + so, ) to

48


In summary, the heterogeneous solid catalyzed gas phase oxida-
tion studies carried out to date are of very limited usefulness
to plume chemistry calculations. First, the literature is essen-
tially void of reaction rate and/or mechanism studies relating
(dry) heterogeneously catalyzed SO, oxidation systems to atmos-
pheric processes. As pointed out in the previous section on
aqueous phase oxidation systems, the most effective catalysis

(88)

occurs in solution. The studies of Corn and Cheng and

(90)

Urone raise an interesting point on the lack of reactivity

of dry particulate vanadium. Because of the recognized activity

(97,98) 54 soften stated that high sul-

of V,0, as an acid catalyst,
fate levels in some areas may be attributable to the use of
vanadium-containing fuels. It would appear from these ambient
temperature investigations, however, that if S03 is produced in
the combustion process, or at least near the stack outlet, the
vanadium-containing particulate matter is not responsible for

excessive sulfate formation in plumes or in the atmosphere in

general.

76


°
@
T

Fraction as H2S0,(g)
° °
b a
T T

O.2-

0.0

| n i L |

“100

200 300 400
Temperature ,C

FIGURE 1. EQUILIBRIUM CONVERSION

S03 (g) + HO (g) == #80, (g) (19)

11

500



that the uncatalyzed reaction produced a negligible amount of
sulfate, and that it was not photosensitive in normal sunlight.
In the presence of a catalyst (FeCl,) the sulfate formation was
found to reach a limiting value after a period of time (typi-
cally one to three hours), and the final sulfate concentration
depended linearly on the initial SO, concentration for a given
catalyst concentration. Furthermore, it was found that the pH
of the solution dropped during the course of the reaction, and
that very little additional sulfate was formed once the pH

reached 2.2.

Junge and Ryan also did a preliminary study to determine the
effectiveness of various salts (at the same concentration by
weight) as oxidation catalysts, and found manganese to be most

efficient (see Table 12). This study was not.meant to be ex-

haustive, however, and it should be kept in mind that in the
atmosphere, the concentrations of various catalysts are not
equal, and thus a less efficient (but more abundant) catalyst
may actually be more important in sulfate production under true

atmospheric conditions.

*
AMMONIA CATALYSIS

Junge and Ryan investigated the enhanced effectiveness of metal

catalysts in the presence of NH.,, and concluded that the NHS

*

It may be argued that promotion of the oxidation of SO, by
ammonia is not catalysis in the true sense of the word; “never-
theless, the term will be used for convenience in the following
discussion with a less-than-formal definition.

52


10
il
12

13

14

TABLES

Equilibrium Constants for the Reaction

803 (g) + 1/2 25 (g) F 83g)

Summary of Studies of SOj Oxidation in Actual
Plumes

SO, Oxidation Studies - Colbert Steam Power
Plant Plume

Pseudo First-Order Rate Constants for Sulfur
in the Los Angeles Basin

Summary of SO, Oxidation Rates in Plumes
Estimated Rates of Theoretically Possible
Homogeneous Removal Paths for SO, ina
Simulated Polluted Atmosphere

SO, Lifetimes for Conversion to H,S0, in
TrOposphere

Estimated Rates of SO, Removal, T = 300°K

Temperature Dependence of Equilibrium Constants
for the SO0,-H,0 System

Summary of Aqueous Phase SO, Mechanisms
Effect of Acidity on Reaction Velocity

Effect of Various Solutions on Sulfate
Formation

Nomenclature and Values Used for Plume SO,
Oxidation Calculations

Summary of SO, Reactions on Various Particulate
Species

iii

14

16

27

28

32

39

41
44

45
49
53

60

73


will be brought out later, this drop-off in catalytic conversion
efficiency may raise questions about the significance of hetero-

geneous gas-phase reactions in plumes.

It has generally been observed in fuel-burning processes that l-
2 percent of the sulfur in the fuel leaves the stacks as S03,
and both heterogeneous and homogeneous reactions have been ‘sug-
gested to account for this result. Expressions developed by
Calderbank 85) and by Kodles (86) | respectively, indicate the
rate of SO, oxidation to be negligible (< 3 ppm So, per hour) at
300°K with 100 ppm S05:

moles SO, converted 31,000 Po, P80,
r, = exp ( - + 12.07) vy
g catalyst-sec RT Pso
2
_ 6 _ 32,414 wy (1/2
r= 1.76 x 10° exp ( ir N50, No, Nso,)

Here the P's are in atmospheres, the N's are mole fractions, T is

the absolute temperature (K) and R isthe gas constant. This low
rate tends to support earlier work (86) which shows that the conversion

of SO, to SO, can occur homogeneously in the flame front (the reactive

3
region of the flame) by stationary state processes:

So, +O+M7S0,+M

2

so, + 0 = SO, + 0, .

3
A recent study by Novakov, Chang, and Harker (87) presents evidence
for the possible role of carbon (soot) particles as catalysts

for the oxidation of SO, in the atmosphere. This evidence is

69


an oy 1300 EST 23 NOV.,1S73 Opt]. 450-1800 EST 22 uuMe, J974. |
* Vf
20
0 Le KM OUT
t0. 100 5
= Awe “
4
20}
4.8 KM OUT
o/

3

M2 KM OUT

TU

KILOMETERS (APPROX.

FIGURE 4. CROSSPLUME 0, AND SQ; VARIATIONS FOR VARIOUS
DOWNWIND DIStancEs@L)

25


SECTION VII

HETEROGENEOUS STUDIES

Considerable attention has been given to the use of catalysts,
especially platinum catalysts, for converting SO, to S03 for
sulfuric acid production. Most recently, platinum and some of
the rare earth oxides have been considered further in their
roles as automotive oxidative catalysts. In this instance, as
in the oxidation that occurs in plumes, the motivation was not
a search for improved catalyst efficiency, but rather to under-
stand the role of the catalyst. A review of the extensive lit-
erature concerning the heterogeneous catalysis of SO, oxidation
for sulfuric acid production (83) would not be generally relevent
to the scope of this report since plume SO concentrations are
orders of magnitude below the levels used in acid processing;

the kinetic concepts involved, however, should be comparable.

(84)

Figure 6 taken from Boreskov presents a brief overview of
the effectiveness of a number of catalysts for oxidizing SO, to
S03 in sulfuric acid production. A most pertinent aspect of
Boreskov's efficiency curves is the sharp drop in the conversion
efficiency of the metal oxides at temperatures below 500°C. The

catalysts most pertinent to plume oxidation shown here might be

V205 and Fe,03 (at much lower temperatures, of course) but as

67


TABLE 13. NOMENCLATURE AND VALUES PSED FOR PLUME SOo
OXIDATION CALCULATIONS © )

Nomenclature Value
General W Effluent dust burden, B/legg 2x 10°?
G Effluent SO, content, ppm 2*5 x 103
D Effluent dilution factor 1073
fy Fraction of total sulphur oxidized 107}
s Droplet suplhate concentrations, mol/l, 0 1
2
Manganese My Oxide molecular weight, g/mol 2:29 x 107
Qin,0,)
a Number of catalytic ions per molecule 3
£, Fraction by weight of dust soluble 2x 10
Iron M as above 1-60 x 10?
(Fe,0,) i
23
n, as above 2
i
fy as above 1072
—= <==

60


order rate constants for the loss of SOs varied from 2 x 107?

to 13 x 1073 min] of 1-6 hours).

(£1 72

(15)

Stephens and McCaldin carried out a series of measurements
in a power plant plume by assuming that particulate matter in
the submicron size range constitutes a conservative (non-varying)
tracer, and relating SO, measurements to this quantity. Friend ©”)
has taken exception to such a procedure on the grounds that SO,
can be oxidized to sulfate in particulate form, thus invalidating
the assumption that particulate matter is a conservative tracer.
This point notwithstanding, Figure 3 (taken from their study)

gives some appreciation of the nature of the results. Half-

lives for curves A, B, and C are:

Curve RH tiy2
A 30 - 408% --
B “40 = 558 _ 144 min
Cc 78 - 803 70 min

The results support the generally established concept that hum-
idity exerts a major influence in the oxidation of SO, in the

atmosphere.

weber '19) has made measurements Of SO, decay using CO, as his
conservative tracer. His measurements were not made in an
elevated plume, but at ground-level sampling sites downwind of

a power plant. Using calculated CO,/S05 ratios for the effluent

18


»k * k

deactivation processes SO, 3? O+ 80, and SO, +M 3 So, +M,
k

the rate constant k, for overall reaction SO, +O+M +4 SO, +M
is given by the expression

ky = ky/ (k,/k3 +M) .

*
(The intermediate So, is assumed to be at its steady state con-
centration.) The ratio ky /k3 was found to be 0.077 mole et,
making k, = 1.4 x 10° 2% mote”? sec? under the experimental

conditions where [M]=10TM mole ete Torr). At atmospheric pressure,
M makes only a slightly more significant contribution (assuming
the mechanism is still valid), leading to a value of k, = 9.3 x

9 2? mole”? sec. If we assume some 104 oxygen atoms/cm? ina

10
plume, the half-life for the reaction is 1.1 x 108 sec (3.5 years),

implying that the process is negligibly slow.

Direct oxidation of SO, by ozone in the gas phase is also a slow
process. In a power plant plume the process is probably of even
less importance because of the rapid removal of ozone by the NO

in the plume. Cox and penkett (22) have considered the oxidation
of SO, in a system containing olefin and ozone. They postulate
that the SO, reacts with an intermediate product (I) resulting
from the reaction between ozone and olefin. The rate of S03 for-
Mation is then expressed by so, = k(SO,) (1). The consequences

of these reactions may be greater in terms of long-range SO, trans-
port than in short-term plume chemistry due to the aforementioned

rapid reaction between O3 and NO.

35


served to neutralize the sulfate formed in the reaction. Its
role in maintaining a high pH and thus a high sulfite concen-
tration is not mentioned, and this is probably a more important
function. A theoretical investigation of the process in fog
droplets estimated 2.9 ug/m3 of sulfate would be formed in an
atmosphere initially containing 20 g/m? of SO, and 3 g/m? of
NH, ("clean country air") with a liquid water content of 0.1
g/m. The same fog in air polluted with 500 ug/m? of SO, and
10 ug/m? of NH, initially should form 26.2 ug/m? of SO, 4 or
nearly an order of magnitude more. No calculations were

attempted for falling drops because of uncertainties regarding

concentration behavior of the catalysts in such systems.

Experimental work on the formation of ammonium sulfate in water
(52)

droplets was done by van den Heuvel and Mason Using water
droplets suspended on a fiber grid and exposed to an airstream
containing controlled amounts of So, and NH3, they found that
the mass of salt (assumed to be ammonium sulfate) produced in
the droplet is proportional to the surface area of the drop and
to time. Because of certain problems with the technique, how-
ever, their results are only semi-quantitative. Furthermore,
application of their results to the atmosphere requires extra-
polation to much lower gas phase SO, and NH3 concentrations than
were used in the study. Allowing this extrapolation, their pre-

dicted rate dependence leads to a rather large conversion rate

of about 2.5%/min in an atmosphere containing 100 ug $0,/m> and

54


Small deviations from standard isotopic sulfur ratios are expres-
sed by a ‘del value':

$s = eae (standard)

- 1] x 1000
325/345 (sample) |

If isotopic equilibrium is attained between SO, and SO, so that

34 32

+ so 32

so, + 3480

so 2 3 '

K
2 30+

the equilibrium constant can be written in terms of 680, and 6sOz

1000 + 650,

K = [000 + 550,

The fraction of 805 converted to S03, 1-f£, is calculated from the

expression

$505 = $80, (fuel) - a(1-£)1000 ,

where a = K-l. This equation is derived from the definition of
K by considering the complete combustion of sulfur in the fuel

and an isotopic material balance.

Surprisingly, the oxidation rate in the plumes of the oil-fired
plant was higher than that in the coal-fired plant:

oil-fired tiy2 = 10 hours

coal-fired ~ 0% loss.
Newman et al. explained that the low oxidation rate in the coal-
fired plume was limited by the low particulate content of the.
plume, where loadings were of the order 10° g/m. No loadings

were reported for the oil-fired plumes.

21


71.

72.

73.

74.

77.

78.

79.

80.

8l.

Volfkovich, S. I., and A. P. Belopolskii, Oxidation of
Sulfites. J. Applied Chem. (U.S.S.R.). 5: 509 (1932).

Haber, F. and O. H. Wansbrough-Jones, Autoxidation (VI).
Action of Light on Sulfite Solutions in Absence and
Presence of Oxygen. 2Z . Physik. Chem. B18: 103 (1923).

Briner, E., and H. Bierdermann, The Role of Ozone as
Catalyst of Oxidation (V). The Ozonization of Sodium
Sulfite--The Influence of the Dilution of Ozone and the
Concentration of Hydrogen Ion. Helv. Chim. Acta.

16: 548 (1933).

Reinders, W., and P. Dingemans, Speed of Oxidation of
Metol by Air and the Effect of Sodium Sulfite. Rec. Trav.
Chim. 53: 239 (1934).

Sillen, L. G., Stability Constants of Metal-Ion Complexes,
Section I. Chem. Soc. Special Publ. No. 17 (1964).

Bassett, H., and A. J. Henry, Formation of Dithionate by
the Oxidation of Sulfurous Acid and Sulfites. J. Chem.
Soc. 1935: 914.

Dana, M. T., D. R. Drewes, D. W. Glover, and J. M. Hales,
Precipitation Scavenging of Fossil-Fuel Effluents: Power
Plant Measurements and Development of a Chemical Kinetics
Scavenging Model. Final Report to EPA, Research Triangle
Park, N.C. (1975).

Johnstone, H. F., and D. R. Coughanowr, Absorption of

Sulfur Dioxide from Air. Ind. Eng. Chem. 50: 1169 (1958).

Nytzell-de Wilde, F. G., and L. Taverner, Experiments
Relating to the Possible Production of an Oxidizing Acid
Leach Liquor by Auto~Oxidation for the Extraction of
Uranium. Proc. 2nd U.N. Int. Conf. Peaceful Uses at
Energy, 2nd, Geneva. 3: 303 (1958).

Bjerrum, J., Stability Constants cf Metal Ion Corplexes
(Supplement). Chem. Soc. Special Publ. No. 25 (1971).

Freiberg, J., Effects of Relative Humidity and Temperature
on Ion-Catalyzed Oxidation of SOj in Atmospheric Aerosols.
Env. Sci. Tech. 8: 731-734 (1974).

86


SECTION VI
HOMOGENEOUS STUDIES IN THE AQUEOUS PHASE

A consideration of primary importance to the question of aque-
ous-phase SO, oxidation is the dissociation behavior of this
material in water. The fact that SO, does form a hydrate which
dissociates in water to yield an acidic solution has been recog-
nized for many years, but until relatively recently the rate of
this dissociation process and its products have been open to spec-
ulation. A great many authors, for instance, have presumed that

undissociated sulfurous acid, H,So , is formed upon dissolution

(45)

of SO, in water. Spectroscopic studies , however, indicate

that, while H,S03 does not exist to any appreciable extent in the

SO.-H,O system, such species as undissociated S05, HSO so,”

272 3° 3°’
pyrosulfite ions (eg. HS,0.°), and hydrates (SO, *nH,O) do exist
in measurable quantities. Furthermore, S05, HSO,”, and SO,” are
the most abundant species, and for most practical applications

one can represent the dissociation process in terms of the

equations
Ko
S09 (g) z 809 (aq) (1)
Ky . +
50> (aq) + H,0 z HSO, +H (2)
K
- “2 = +
HSO, 3° SO, +H : (3)


ll.

12.

13.

14.

15.

16.

17.

18.

19.

20.

2l.

Verhoff, F. H., and J. T. Banchero, Predicting Dew Points
of Flue Gases. Chem. Eng. Prog. 70: 71-72 (1974).

Gartrell, F. E., F. W. Thomas, and S. B. Carpenter, Atmos-
pheric Oxidation of SOj in Coal-Burning Power Plant Plumes.
Am. Indust. Hygiene J. 24: 113-120 (1963).

Dennis, R., C. E. Billings, F. A. Record, P. Warneck, and
M. L. Arin, Measurements of Sulfur Dioxide Losses from
Stack Plumes. APCA 69th Annual Meeting, New York City
(June, 1969).

Coutant, R. W., E. L. Merryman, R. E. Barrett, R. D.
Giammar, and A. Levy, A Study of the Fate of SO, in Flue
Gas. Battelle's Columbus Laboratories Final Refort to
EPA/API/BCR, Inc./EEI (March, 1972) 12 p.

Stephens, N. T., and R. O. McCaldin, Attenuation of Power
Station Plumes as Determined by Instrumented Aircraft.
Env. Sci. Tech. 5: 615-621 (1971).

Weber, E., Contribution to the Residence Time of Sulfur
Dioxide in a Polluted Atmosphere. J. Geophys. Res.
75: 2909-2915 (1970).

Newman, L., J. Forrest, and B. Manowitz, The Application of
an Isotopic Ratio Technique to a Study of the Atmospheric
Oxidation of Sulfur Dioxide in the Plume from an Oil-Fired
Power Plant. Atmospheric Environment. 9: 959-968 (1975).

Newman, L., J. Forrest, and B. Manowitz, The Application of
an Isotopic Ratio Technique to a Study of the Atmospheric
Oxidation of Sulfur Dioxide in the Plume from a Coal-Fired
Power Plant. Atmospheric Environment. 9: 969-977 (1975).

Newman, L., M. Smith, J. Forrest, W. Tucker, and B. Manowitz,
A Tracer Method Using Stable Isotopes of Sulfur. Proc. Amer.
Nucl. Soc. Topical Mtg. on Nucl. Methods in Env. Res.,
Columbia, MO.

University of Utah Research Institute, Rate of Conversion
of Sulfur Dioxide in a Power Plant Plume to Particulate
Sulfate. Final Report, Phase I, APS Contract No. 75-00634,
funded by WEST Associated (May, 1975).

Davis, D. D., and G. Klauber, Atmospheric Gas Phase Oxida-
tion Mechanisms for the Molecule SO,. Int. J. Chem. Kin,
Symp. No. 1 (1975), Proceedings of the Symp. on Chemical
Kinetics Data for the Upper and Lower Atmosphere, 13 p.

81


rates by manganese and iron catalysts. In the case of iron

catalysis, some of the assumptions made are based on the work

(79)

of Nytzell-de Wilde and Taverner The derivation led to

these rate expressions:

Rate of S05 oxida-

tion by Mn“catalyst = 2

2
22.4 Kiciv
6

x 10°% per min.

107 GD

Rate of SO, oxida-

'
; DWK,.(1 - £_)K,n.f
tion by Fe'catalyst _ 5 54 x 103 H of iii

+ % per min.
[H"] M;

where the symbols are defined in Table 1l3,as are typical values
for each parameter. Using these values and assuming values for
tH), Ky and Ki Foster estimates SO, rates of 0.09%/min for Mn
and 0.15-1.5%/min for Fe, suggesting that iron oxides are the
most important catalysts for aqueous phase so, oxidation in
plumes.

(57) considered the kinetics of the oxidation

Matteson et al.
mechanism using a manganese sulfate aerosol catalyst. A theore-
tical analysis of the mechanism (see Table 10) leads to the con-
clusion that the rate is proportional to the square of the initial
aqueous nit concentration (Foster had assumed a similar depend-
ence in his calculation of the manganese-catalyzed rate). The
rate constant was found experimentally to be 2.4 x 10> wist,
It was also found that almost no sulfate is formed when the re-

lative humidity is less than 95%, probably because the aerosol

59


22.

23.

24.

25.

26.

27.

28.

29.

30.

31.

32.

Friend, J. P., Attenuation of Power Station Plumes as
Determined by Instrumented Aircraft. Env. Sci. Tech.
6: 172-173 (1972).

Whitby, K. T., B. C. Cantrell, R. B. Husar, N. V. Gillani,
J. A. Anderson, D. L. Blumenthal, and W. E. Wilson, Jr.,
Aerosol Formation in a Coal-Fired Power Plant Plume.
Presented before the Division of Environmental Chemistry,
American Chemical Society, New York (April 1976).

Roberts, P. T., and S. K. Friedlander, Conversion of SO
to Sulfur Particulate in the Los Angeles Atmosphere.
Env. Health Perspectives. 10: 103-108 (1975).

Calvert, J. G., Interactions of Air Pollutants--X. The
Effects of SO, Addition of the Reactions in the Simulated
Polluted Atmosphere. National Academy of Sciences, Con-
ference on Health Effects of Air Pollutants, Washington,
D. C. (October 3-5, 1973).

Castleman, A. W., R. E. Davis, H. R. Munkelwitz, I. N.
Tank, and W. P. Wood, Kinetics of Association Reactions
Pertaining to H,SO4, Aerosol Formation. Int. J. Chem. Kin.
Symp. No. 1 (1995) . Proceedings of the Symp. on Chemical
Kinetics Data for the Upper and Lower Atmosphere, 629-640.

Gerhard, E. R., and H. F. Johnstone, Photochemical Oxida-
tion of Sulfur Dioxide in Air. Ind. & Eng. Chem. 47:
972-976 (1955).

Sidebottom, H. W., et al, Photooxidation of Sulfur Dioxide.
Env. Sci. Tech. 6: 72-79 (1972).

Leifer, R., J. P. Friend, and M. Trichon, The Mechanism of
Formation of Stratospheric Aerosols. J. Atmos. Sci.
30: (1973).

Cox, R. A., Particle Formation from Homogeneous Reactions
of Sulfur Dioxide and Nitrogen Dioxide. Tellus. 26:
235-240 (1974).

Jaffee, S., and F. S. Klein, Photolysis of NOz in the
Presence of SO7 at 3660A. Trans. Far. Soc. 62: 2150-
2157 (1966).

Cox, R. A., and S. A. Penkett, Aerosol Formation from
Sulfur Dioxide in the Presence of Ozone and Olefinic
Hydrocarbons. J. Chem. Soc. Far. Trans. I, Part 9.
68: 1735-1753 (1972).

82


studies and the need for considerably more data. The increase
of SO, oxidation rate with increasing relative humidity is
fairly well accepted; on the other hand, the results do not
provide overwhelming support of this conclusion. Similarly,
the dependence of SO, oxidation rates on heterogeneous cataly-
sis is accepted by many, yet none of the plume experiments has
attempted to link these variables. In summary then, one can
only say at this point that SO, oxidation rates in plumes may
vary from essentially zero to the order of 50 percent per hour.

A clearer description of actual SO, oxidation processes in

plumes will have to await considerably more experimentation.

29


SECTION II
INTRODUCTION
OBJECTIVE AND SCOPE

This report reviews the scientific literature relating to the
oxication of sulfur dioxide in power plant plumes, and provides
recommendations on conversion rates to be applied in atmospheric
dispersion models, based on current knowledge in this field.

The importance of these two objectives derives both from the
presently confused state of our knowledge regarding atmospheric
oP) oxidation processes and from the potentially great problems

associated with our rapidly increasing sulfur dioxide emissions.

Recognizing these problems, a program has been initiated to pro-
vide a more complete means for evaluation of the atmospheric im-
pact of increased fossil fuel combustion. The first component

of this program has been a review of the behavior of plumes emit-
ted from large, elevated (primarily fossil fuel combustion) sour-
ces, which has been documented in the report Tall Stacks and the

Atmospheric Environment ‘1) , From this report it is highly evi-

dent that our present ability to predict the atmospheric impact
of increased fuel utilization is limited by lack of knowledge
regarding in-plume SO, conversion processes. The present report

is intended to help rectify this situation somewhat by providing


Wang and Himmelblau were apparently unaware, however, that Eigen
et at, (47) had performed a previous definitive measurement of the
rates of ionization of hydrated S05, which strongly indicates their
later results to be in error. Eigen's measurements indicate that
ionization approaches completion at times of the order of 1075

seconds, thus enabling these reactions to be regarded as essentially

instantaneous for this assessment of oxidation processes.

TABLE 9. TEMPERATURE DEPENDENCE OF EQUILIBRIUM
CONSTANTS FOR THE S02-H20 system (48)

Temperature Ky Ry Ky
25°C 1.24 M atm -0174 M 6.3 x 1078 M
15 1.83 0219 7.9 x 107°
10 2.25 0247 8.9 x 1078
0 3.46 .0319 «11.4 x 1078
-3 3.97 10346 © 12.4 x 107%

The literature pertaining to the aqueous phase homogeneous ox-
idation of SO, is in some ways more pleasing than that for the
gas phase (a summary is given in Table 10). For instance, there
is general agreement that in the aqueous phase, the oxidation of
S(IV) by 0 must be catalyzed in order to proceed at an appreciable

rate, and almost without exception the suggested catalyst has

44


It cannot be overemphasized that the oxidation of SO, in the
atmosphere is a highly complex process, and that the rate is and
should be highly dependent on the nature of the plume and the
existing meteorological conditions. Because of the extreme vari-
ability of data pertaining to atmospheric $0, oxidation processes,
it is difficult to provide any meaningful recommendation on rate
expressions to be used for practical modeling purposes. It can

be stated with some certainty, however, that insofar as modeling

for scoping purposes is concerned, the results available thus far

do not justify the application of rate expressions any more elab-
a (S05)
dat

orate than the psuedo~first order form

= -k (S05).

Since SO, oxidation processes are typically highly complex, they
will not be expected to adhere to the above linearized expression
accurately over extensive ranges of conditions. The basically
unsatisfactory state of this field, however, leaves us with lit-
tle reasonable choice but to proceed with this approach. Recog-
nizing these factors, and with a warning against indiscriminate
use, it is suggested that, for general purposes of model calcu-
lations, first order SO, oxidation rates in the range 0.5 to
10%/hr are consistent with the bulk of the literature, and a
reasonable value for many situations is 2%/hr. In the investi-
gation of specific mechanisms, however, it is suggested that
rate data for the specific reactions involved be used whenever

possible.


Davis (21) has also reviewed some aspects of homogeneous So,
oxidation kinetics with a slightly different approach from that
of Calvert. Davis considers the following processes and esti-

*
mates SO, lifetimes for the lower troposphere as shown in

Table 7:
(1a) $0, + hv (2400-3400 a) > 805
(1b) S05 + 0, + (S04)
(2) SO, +O +M->S0,+M
(3) 80, + 0,(*A) + (804)
(4) $0, + 0, + 80, + 0,
(5) SO, + NO, + SO, + NO,
(6) SO, + N30, > So, + N30,
(7) 80, + HO, + SO, + OH
(8) SO, + OH +M+ HSO, + M
(9) SO, + CH,0, > SO, + CH,O

Reaction 8 is the most probable for the homogeneous oxidation of
SO, in the atmosphere. Thus, the structure and reactions of the
product of reaction 8, HSO3, is very critical to acid aerosol
formation in the atmosphere. It is on this basis that Davis
puts forth the extensive series of reactions on the oxidation of

HSO, discussed in the section on plume studies.

*The term "lifetimes" is frequently seen in the literature, and
is equivalent to "half-life".

38


A few years after the TVA Study, a study was carried out by Den-

nis et al. of the GCA Corporation ‘23)

, in which SO, losses in
stack plumes were determined downwind from a 330-MW coal-fired
power plant. Since it was not feasible to measure SO directly,
SO losses were determined relative to SFe, a commonly-used non-
reactive tracer. The results of this program were inconclusive.
The half-life of SO, under the field conditions ranged from 1.0

to 2.8 hours with a mean value of 1.5 hours. There were some dis-
crepencies in this study in extrapolating the field SO02/SF, ratios
back to the expected source value for this ratio: Some extrapola-

tions suggested fairly rapid half-lives, such as 0.13 hours. Rela-

tive humidities in this program ranged from 36-53 percent.

As a follow-up to this study a laboratory program was carried out
by Coutant et al., at Battelle-Columbus Laboratories primarily
focused on determining if SO, decay was as rapid as the GCA data
suggested within a half mile of the power plant 24), In this
program, a time-temperature simulation of the stack gases leaving
the furnace was established. The burner was fueled with three
types of coal (at 30 lb/hr) and two oils (at 3 gal/hr) to assess
the variation of the SO oxidation rate with fuel composition.
When burning coal, there was about a 10 percent loss of SO, in

the flue gas within the boiler-economizer section, while there

was essentially no loss of SO, when burning oil. In the plume

region, loss of SO was very dependent on humidity. First-

17


100

TTT TT TT

PPM H2SOzq in Gas Phase
°

Lisle and Sensenbaugh

0.!
——— Verhoff and Banchero
(for 12 % water)
0.01 ! i | 4 L
90 100 Ho 120 130 140 150

Temperature,C

FIGURE 2. DEW POINT OF H,S0, (g) (20)

12

160


SECTION I

CONCLUSIONS

The available literature concerning the oxidation of SO, in plumes
does not allow firm quantitative conclusions to be drawn concern-
ing absolute rates or mechanisms. Reported rates cover a wide
range, and few details of the relevant mechanisms are well estab-
lished. Despite these uncertainties, however, one can state with
some confidence several characteristics of in-plume SO conversion
processes. It is known, for example, that many factors strongly
influence the rate of oxidation of SO,- The evidence collected
thus far indicates that the gas phase oxidation rate increases
with increasing relative humidity. The concentrations of catalysts
(in the case of heterogeneous oxidation) and reactive species (for
homogeneous oxidation) also exert a profound influence on the rate.
In aqueous systems, pH, temperature, and catalyst concentration are
perhaps the most significant factors in determining the oxidation
rate. Furthermore, several complex mechanisms may be involved in
the oxidation under various conditions. Indeed, considering the
magnitude of the problem, it would be extremely fortuitous if a
single rate expression could be applied under all meteorological
conditions and all ambient atmospheres for all plumes generated

by all types and grades of fossil fuels. This, however, is little
comfort to the modeller, whose job it is to reduce the complex
system to a managable number of variables, relate these variables
with a sound mathematical formalism and perform calculations to
yield results which conform qualitatively and (hopefully) quanti-

tatively to the actual situation existing in the atmosphere.


and catalyst concentration, which leads to the expectation that

the rate really is highly variable. Thus, the conclusions which

do emerge from the works reported are mainly qualitative:

(1)

(2)

(3)

(4)

(5)

(6)

Under favorable meteorological conditions, aqueous
phase oxidation of SO, must be considered, and in some
cases it may be the predominant mechanism.

Power plant plumes are significant sources of the met-
als (especially manganese and iron) which are able to
catalyze the oxidation.

The influence of ammonia on the reaction is of parti-
cular importance, both in the formation of ammonium
sulfate and in its ability to increase the reaction
rate by maintaining a high pH.

Ozone is a potentially important aqueous phase oxidizing
agent for S05, although its importance inthe early plume
may be doubtful because of depletion by reaction with NO.
Aerosols of MnSO, and other species may be important
as sources of catalysts, as well as reaction sites
when the aerosols are hydrated.

Because of the highly sensitive nature of the reaction
(or reactions) to existing ambient conditions, typical
rates probably range from about 0.1 to 2.0%/min, al-

though even higher rates may be possible.

66


The authors speculate that a catalytic mechanism may be the domi-
nant pathway for oxidizing SO, in plumes and refer to indigenous
vanadium as a potential catalyst. This conjecture is not to be
discounted, but in the same vein, one might also expect sufficient
Fe,0, in the particulate matter from the coal-fired plant to cata-
lyze the oxidation of SO, and there is growing doubt concerning
the catalytic efficiency of vanadium in plumes. (The reader is
referred to the section on heterogeneous oxidation of So, for other

comments on vanadium catalysis.)

The BNL investigators attempted to fit the decay to first and
second order kinetics, and based on a second order mechanism

lpr “+, which accounts for the

arrive at a rate constant of 1 ppm”
10 hour half-life referred to earlier. No specific dependence
on humidity was noted in this analysis, although the RH levels
varied from 40-95 percent. In fitting their data to a second
order mechanism the authors do not mean to suggest that the oxi-
dation process arises from a bimolecular reaction of S05, but
that the process involves the reaction of SO, with active part-
icles in the plume through water associated with these particles.

It is their opinion that any homogeneous mechanism is inconsis-

tent with the isotope ratio data.

A study carried out by the University of Utah Research Instit-
ute (20) at the Four Corners power plant in New Mexico in 1974,
also provides evidence for low SO, oxidation rates in coal-

fired power plant plumes. This study was part of an on-going

22


(2), Kellog, et al. (3), Cadle (4), M. Bufalini (5), A. P.
Altshuller and J. Bufalini (6), Harrison, et al. (7), Pierrard
(8)). In general, the available literature pertaining to each
of the topics to be discussed in this report varies widely in
both quantity and quality, making the extraction of relevant
data and conclusions difficult. The aim of this review, then,
is to extract the data which are available from the literature,
however qualified they may be, and also to draw attention to

the areas most lacking in substantive conclusions.


Table 4. Pseudo First-Order Rate Constants for the
Reaction SO,+1/2 02 k SO3 in the Los Angeles
Basin (24).

Date Time of arrival Starting
(1973) at Pasadena k, %/hr location
(PST)

July 10 1300
1400
1500

1 El Segundo
3 " "
9
1600 13.
2
8
8

" "
” "

July 25 1400 12.
1500
1600

8 Alamitos Bay
2 El Segundo
8 0 "

July 26 1200 4, Alamitos Bay
1300 5. " "
1400 7.

4,

3

6

6 " "
1500 7

27


requirements. The units of ky and kL, and thus M and N, are
(concentration) 27" et. The reaction order r is the total number
(possibly fractional) of atoms or molecules; in the general case
described here, r=2 for both the forward and reverse reactions.
If, however, reactant A is present in such great quantities com-

pared to B that its concentration remains virtually constant, A

may be combined with ke so that
(A[B]/dt) , = -k,IA][B] & ~k,[B] :

Here, although the reaction is actually second order, a pseudo-

1
first order rate expression with constant ky can be applied.
Such approximations are often valid in the atmosphere, especially

when a trace substance is oxidized by oxygen, a major constituent.

The half-life of a species is the time required for half of it to
react. In the special case of a first order (or pseudo-first
order) reaction, this time is simply (ln 2)/k, where k is the rate
constant. This quantity is often used to compare speeds of com-

peting reactions.

If the reaction is allowed to proceed undisturbed for a long
enough time, a state of dynamic equilibrium will be established
when the rate of the forward reaction is exactly equal to the
rate of the reverse reaction. That is, although both reactions
are still proceeding, there are exactly as many molecules of
each species formed by one reaction as are destroyed by the

other. Thus,

94


82.

83.

84.

87.

88.

89.

90.

91.

92.

93.

Jones, P. W., and A. H. Adelman, Photosulfoxidation of
Hydrocarbons in the Liquid Phase. Env. Sci. Tech. 6: 933
(1972).

Duecker, W. W., and J. R. West, The Manufacture of Sulfuric
Acid. A.C.S. Monograph No. 144, Reinhold (1959).

Boreskov, G. K., Mechanism of the Oxidation of Sulfur
Dioxide Gas on Oxide Catalysts. J. Phys. Chem. (U.S.S.R.).
14: 1337-1346 (1940).

Calderbank, P. H., Contact-Process Converter Design. Chem.
Eng. Prog. 49: 585-590 (1953).

Kodles, B., et al, Study of the Oxidation of S02 on a
Vanadium Catalyst in the Region of Internal Diffusion.
IV Int. Congress on Catalysis, Proc. Symposium III,
2477-2499 (1968).

Novakov, T., S. G. Chang, and A. B. Harker, Sulfates as
Pollution Particulates: Catalytic Formation on Carbon

(Soot) Particles. Science. 186: 259-261 (October 18,

1974).

Corn, M., and R. T. Cheng, Interactions of Sulfur Dioxide
with Insoluble Suspended Particulate Matter. APCA Journal.
22: 871-876 (1972).

. Chun, K. C., and J. E. Quon, Capacity of Ferric Oxide

Particles to Oxidize Sulfur Dioxide in Air. Env. Sci.
Tech. 7: 532-538 (1973).

Urone, P., et al, Static Studies of Sulfur Dioxide
Reactions in Air. Env. Sci. Tech. 2: 611-618 (1968).

Smith, B. M., J. Wagman, and B. R. Fish, Interaction of
Airborne Particles with Gases. Env. Sci. Tech. 3: 558
(1969). a

Low, M.J.D., A. J. Goodsel, and N. Takezawa, Reactions

of Gaseous Pollutants with Solids (I). Infrared Study
of the Sorption of SOz on CaO. Env. Sci. Tech. 5: 1191-
1195 (1971).

Goodsel, A. J., M.J.D. Low, and N. Takezawa, Reactions of
Gaseous Pollutants with Solids (II). Infrared Study of
Sorption of SOz on MgO. Env. Sci. Tech. 6: 268-273
(1972).

87


SECTION III
THERMOCHEMISTRY OF SO, OXIDATION

In the gas phase, the production and subsequent reaction of sul-
furic acid (H5S0 4) in a plume involves the oxidation of SO, to
S03, hydration of S03, condensation of the resulting acid, and
reactions with particulate matter. Equilibrium constants for

(9)

the oxidation of SO, are presented in Table l. It becomes

apparent

Table 1. Equilibrium Constants for the Reaction

> (9)
$09 (g) + V2 %(g) * $03 (9) :
Temperature Equilibrium constant (Peo P55 ei’)
OK °0 atm71 3 2 2
300 80 2.069 x 10%?
400 260 1.088 x 108
500 440 2.608 x 105
x 103

600 620 4.892

that, under conditions pertinent to gases leaving the stack
(<500 °K) and entering the plume (300-400 °K), the equilibrium
so heavily favors the formation of S03 that in the absence of
kinetic factors, there would essentially be no SO, present.
These kinetic factors, and the ways in which they influence

the oxidation rate, will be addressed in subsequent sections

of this report.


dependent on the amount of material present. For example, in

the reaction A+ B+C+D, A and B are converted, and
(a[A]/dt) - = (4[B] /dt) - = -k, [A] [B]

where the brackets denote concentrations. Similarly, A and B

are both produced by the reverse reaction, (C + D+ A+B), so
(a[Al/dt) . = (4[B]/dt) . = +k[C) [D]
and the overall rate of change of A and B is
a[A]/dt = a[B}/dt = k,[C][D] - k, [A] [5] .
In the same way,
a(Cj/dt = a[D]/dt = k, [A] [B] - k, [Cl] [D] = -d[A}/dt .

The rate constants Ke and kL can generally be expressed in the
Arrhenius form:
ke = M exp(- E,/RT)

ky

N exp (- E,/RT)

where M and N are constants with only small temperature depend-
ences, Ee and EL are the molar activation energies for the for-
ward and reverse reactions, respectively, and R and T are the gas
constant (1.98 cal/mole-°K) and the absolute temperature. The
exponential term reflects the kinetic energy distribution of the
reacting molecules, that is, the probability that collisions will

be sufficiently energetic to overcome the activation energy

93


12.

13.

14.

15.

16.

17.

18.

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to account for these features is a free radical chain adapted

from Backstrom (70) ,

= k
=. 3+ 51 2+ eco 7
3 + Co(H,0)~" +” Co(H,0) Z* +-s0

so. 6 3

- ky -
80,” + 0, +°+S0,
- = k3 =
"S057 + $0," +7 so," + "so,
= = =
$0," + S0,~ +" 280,
. so, +. $0,” > inert products
. 50,7 + SO,” + inert products
- - ks
* S05 + ° SO, +" inert products.

It has been mentioned before that many of the studies on SO,
oxidation have been done using rather high SO, concentrations,

requiring a large extrapolation in order to apply the results

(63)

to the atmosphere. Brimblecombe and Spedding attempted to

correct this situation by measuring the oxidation of SO, at

concentrations of about 1075

6

M in aqueous solutions with Fe(III)
at about 107 M acting as the catalyst, using a radiochemical

method of analysis. The sensitivity to the iron catalyst, even
at these low SO, concentrations, is indicated by the fact that
even in the purest water obtainable, the oxidation proceeded at
a measureable rate, presumably due to trace iron at concentra-

tions on the order of 1078 M. They suggest that a free radical

62


r

5

8

B 60

a

5 40 V,0,

Fd 28 Ci Oy

2 zw

§ Fe,03 “CyO |
° \ 1 1 fl
300400 500 600 700

Temperature, C

FIGURE 6, CONVERSION OF S07 70 $03 IN THE PRESENCE
OF SEVERAL CATALYSTS (

68


Table 2. SUMMARY

OF STUDIES OF SO, OXIDATION IN ACTUAL PLUMES.

Study

Species Measured

Comments

Gartrell,et al.
(1963) (12)

Dennis,et al.
(1969) (23)

Coutant,e¢
14)

al.
(1972) ¢

Stephens and

McCaldin (1971) 5)

Weber (1970) (16)

TTo5e) (27, 28; 29)

University of
Utah (1975) (20

Davis Klauber
(2975) C21)

Whitby, et al.
(1975) (23)

S05, SO,

SO, relative to SF¢

S05

SO2 relative to sub

micron aerosol
SO, relative to C09

32g, 345

Particulate (NH
S04 relative to
ambient S02

NO, NOz, 03, SO2

42

Aerosol

Highly variable oxidation
rates. Positive correla-
tion with humidity.

SO9 half life from 1.0 to
2.8 hours (mean 1.5 hours).

Laboratory simulation,
various fuels. SO, loss
highly dependent on hum-
idity.

Oxidation rate dependent
on humidity.

Increasing oxidation rate
with increasing humidity.

Higher oxidation rates in
plumes from oil-fired
plants than coal-fired.
No specific humidity dep-
endence noted.

Low oxidation rates, pos-
sibly due to low humidity.

Ozone 'bulge’ noted down-
wind.

Results of single flight
Oxidation rates of 1.5 -
1.8 %/hr inferred.

14


k, {Al [B] - k, [Cj [D] =0 -

In this case we can write an equilibrium expression for the

reaction,
>
A+B+cC+#D

and define an equilibrium constant:

x= £2 [clip
~ Tarsy *

From the preceeding discussion, it is essential to realize that
chemical conversion in macroscopic systems can be limited by

either or both kinetic and thermodynamic factors. The kinetic

factors are essentially determined by the energy available to
the system: if only the most energetic collisions are able to
culminate in reaction because of a high activation energy, the
rate will be slow. Kinetic barriers can be overcome either by
providing more energy to the system (usually as heat) or by ap-
plying an appropriate catalyst to lower the activation energy
of the system. The thermodynamic factors are related to the
equilibrium state of the system, and are somewhat more difficult
to manipulate in situations where they are rate limiting. Be-
cause they only influence the activation energy, and not the
equilibrium composition, catalysts are unable to affect the
thermodynamic factors. The equilibrium constant is a function
of temperature and can be determined from the definitions given

previously:

95


c & 12719768
aa Rm 01/02/69
L © 01/03/69

&
¢
Bi:
Peels ‘
B yok wg 9 tonedhnan
Ee cetie  & oe fe
E rr rn
° of © 68
L
:
pop pe
3) eet
Phe Age, fren

FIGURE 3. SULFUR DIXOIDE DECAY RATES IN THREE RELATIVE
HUMIDITY RaNGEs (15) .

19


In general, it should be reiterated that the aqueous phase
cxidation of $0, is highly dependent on meteorological condi-
tions, since even the fastest mechanism can produce no sulfate
if there is no aqueous phase present. Furthermore, even when
there is an aqueous phase present, scavenging efficiency,
solubility, and diffusion must be considered. For these
reasons, aqueous phase oxidation must be considered somewhat
differently than gas phase oxidation when applied to power

plant plumes.

Although oxidation of sulfite ions and "sulfurous acid" by
oxygen was investigated as long ago as 1g9g 65-74) | much of
this early work cannot be considered quantitatively since,
among other things, there is often no record of purity pre-
cautions applied to regents. These works did show the oxida-
tion to be a long chain reaction, and as such to be quite
sensitive to both positive and negative catalysts (catalysis
and inhibition), but the details of the mechanism are still
uncertain. In his investigation of the effect of inhibitors
on the reaction, Backstrém 70) found that his data fit the
relation

v = A/(B+m)

(where v is the velocity of oxygen absorption, A and B are
constants, and m is the inhibitor concentration) although

deviations were noted for low values of m.

47


degree of completion of the reaction. Moving from left to right,
the initially separated molecules A and B approach and collide,
increasing the potential energy of the system until a maximum is
reached, and then separate forming products C and D. If the
energy of the products is lower than the energy of the reactants,
the first figure is appropriate and the reaction is termed exo-
thermic; otherwise it is endothermic as shown in the second fig-
ure. The amount of energy gained (or lost) by the system during
the course of the reaction is called the enthalpy, denoted AH is
negative for exothermic reactions; reactions of this type corres-
pond to lower energy of reaction products and thus, are said to

be "favored" energetically.

Because of the presence of the maxima, or "energy barriers" in

the figures, all collisions between A and B do not result in
reactions. If a collision is not sufficiently energetic to

reach the top of the energy barrier, the molecules must separate
and remain as A and B; in fact, even with enough energy input,
there is still a chance that the necessary rearrangements will

not take place to allow the formation of products. The height of
the energy barrier is called the activation energy of the reaction,
labled Eg: The action of chemical species called catalysts can
thus be understood: without reacting themselves, they provide a

means whereby E, can be lowered, allowing less energetic colli-
sions to pass over the energy barrier. Figure 7 also reveals

that with sufficient energy, C and D should be able to react to

91


Homogeneous gas phase oxidations proceed via second or third
order processes between SO, and molecular or free radical oxidi-
zing agents. Table 6 presents a compilation of homogeneous
reaction paths drawn up by Calvert (25) , Although the merits of
some of these reaction steps may be questionable, they do show
that significant reaction rates can be expected from homogeneous
reactions in a plume. The compilation is of special interest

in that it offers potential support for rates of 1.7-4.7 percent

per hour.

Since the publication of this compilation, Castleman and the
group at Brookhaven ‘79) have examined the reaction SO, +OE+ HOSO

in some detail. They report a rate constant of 6 x 10733 em?

noleé? sec) (887 ppm?

min7!) for this reaction. If the OH
radical concentration is 107 molecules/em?, SO, conversion rates
of 2 percent per hour would be possible. This conversion rate
is of course very dependent upon the OH concentration in the
plume: if the assumed OH concentration is an order of magnitude

too high (a possibility) then the rate of ~0.23 percent per hour

*
reported in Table 6 would agree with the Castlemen rate data.

In the Castleman study, the reaction SO, + H,O + M+H,SO,+M is

os
A basic dilemma that will become apparent in the following dis-
cussion is the reliability of the rate data relative to the con-
centration data of radical species in the atmosphere. Obviously
this dilemma becomes even more severe in chimney plumes where
the concentrations of radical species are even more in doubt.

31


TVA study, presents a brief summary of the oxidation rates
resulting from measurements made in the plume of the Colbert
power plant, near Wilson Dam, Alabama. The results show @at
in periods of low humidity, SO, oxidation was quite low, rang-
ing from 1-3 percent per hour. At high relative humidity, a
rate of some 30 percent per hour (55% over 108 minutes) was

observed.

Probably the most significant finding coming out of this study
was the extreme variability of SO, oxidation rates. This
variability may have been the result of changing meteorological
conditions, the difficulties of analysis, or factors unknown.
It is of interest to note that in spite of the tendency of
later workers to attempt to relate back to these studies, it
was never the intent of the authors to consider this a defini-
tive study - in the words of the authors, "These limited data
do not provide a basis for an accurate estimate of the absolute
rate of $0 oxidation after emission." The fact remains, how-
ever, that as the first major study of its kind, this work
provided a foundation for subsequent studies by pointing out
some of the important variables and giving at least a quali-

tative idea about the nature of SO, - plume chemistry.

15


The other point which must be kept in mind is the time required
for appreciable oxidation. A fog or mist could be in the vic-
inity of a plume for times on the order of hours. Precipitation
falling through a plume, however, does so in a matter of minutes
or less. Thus, in order to contribute significantly to the sul-
fate formation in the aqueous phase, a mechanism must have a
time constant’ of the same order as the time that the phase is in

contact with the plume.

METAL CATALYSIS

Concurrently with the above work on the SO,-NH,-water system,

there have been several important studies on the effect of metal

catalysts in the oxidation of S0,- Work by Johnstone and Cough-
anowr ‘78) indicated that, for high catalyst concentrations, the
oxidation could be assumed to occur within a spherical shell at
the surface of the drop. This shell becomes thicker as the cat-
alyst concentration decreases, or as the SO, concentration at
the surface increases, and for some critical value of either

parameter, of course, the entire drop is involved.

An excellent review and extension of the earlier theoretical

(56) | whose interest was in the oxida-

work was given by Foster
tion of SO, in power plant plumes. After a lengthy analysis of

droplet (especially MnSO,) growth, he considers the oxidation

*The time constant for a mechanism may be thought of as the
half-life of the rate-limiting reaction.

58


Unfortunately, the development of an entirely general model is
highly impractical. It may, however, be possible to model vari-~
ous specific situations with a degree of success by using avail-
able data with appropriate parameterizations, the nature of

which will depend upon the exact requirements of the model.

The oxidation of S05 has been studied both in the laboratory and
in the field. The results of those studies are divided into

four broad classifications and summarized below.

Plume Studies

1) Measured SO, oxidation rates range from 0 to 55%/hr.
The dependence of the rate on meteorological variables,
especially relative humidity, is recognized, but there
is no way to control these variables in field experi-
ments.

2) Recent studies indicate lower rates than previously
reported, but as yet only one study has measured sul-
fate formation directly.

3) It is recognized that the rate of oxidation varies
spacially in the plume; the results of one study, how-
ever, indicate a high rate near the stack and a lower
rate downwind, while another suggests the opposite
situation.

4) Two recent studies indicate a correlation between ozone

in the plume and the oxidation of S05, with the oxida-


SECTION X

APPENDIX

THERMODYNAMICS AND KINETICS: A REVIEW

Because of the volume of material covered by this report, little
attempt has been made to rigorously define the more common chem-
ical terms or to discuss kinetic and thermodynamic concepts. It
is hoped, however, that this material will be useful to the non-
chemist. This section is intended to review basic kinetics and
thermodynamics for these readers, and to simultaneously define
some of the more general terms and concepts used throughout the
report, for the sake of completeness. A more detailed discussion
of thermochemical kinetics than the necessarily brief overview
presented here may be found in any textbook on physical chemistry.
It is essential to remember that the primary factor. determining
whether molecular or atomic species will react is their energy;

species having lower energy are favored, and all reacting systems
will tend toward configurations having the lowest possible poten-
tial energy. This behavior can be illustrated in terms of the

bimolecular reaction,

AtB+cCH#+oD
whereby reactants A and B are converted to products C and D. Fig-

ure 7 shows how the potential energy of the system varies with a

parameter called the reaction coordinate, which indicates the

89


Smith and vrone ‘3)

have performed studies of the photochemical
oxidation of SO alone and in the presence of NO», propylene
(Ce) and water vapor. The initial SO. concentration throughout
the study was 2 ppm. In air containing only S05, the initial
photochemical rate (aSO,/dt) was 1.74 x 1074 ppm/min (about 0.55%/
hr). The introduction of NO, increased this rate when the S0,:NO>
ratio was 1 or 2, but decreased it when the ratio was 0.6 or less.
When both NO, and propylene were added, the rate increased by some
two orders of magnitude over the rate for the SO,-air system, and
was found to be a function of both the propylene and NO, concen-
trations. The addition of water vapor at 50% relative humidity
to the So, - NO, systemwas found to increase the rate about ten-
fold over that for the dry system, but for unknown reasons a simi-
lar addition to the 50, - NO, - propylene system had no effect.

Wilson and Levy (34135)

have also examined the smog process in
irradiated SO, - NO - hydrocarbon systems. This work was directed
at observing the effects of SO, in smog rather than the reverse.
It was generally observed in this work, however, that the decay

of SO, increased in proportion to the reactivity of the hydro-
carbons. No quantitative values came out of these studies.

In a current program under EPA-sponsorship, D. F. Miller 36) is
studying the homogeneous rate of oxidation of S05 in photochem-

ical smog systems. The studies are being conducted in a large

610-cu-ft smog chamber, and while there are admittedly certain

36


k EL - E
f£_M xr £ M -AH
Ke ge y exp( ag) = F exp (Re .

This indicates that K decreases with increasing temperature for

an exothermic reaction.


(21)

Table 7, SO, LIFETIMES FOR CONVERSION TO H,SO, IN TROPOSPHERE
Second Concentration
Reaction Species Molecules/cm Lifetime days
1 -- -- 4x 107
2 © (3p) ~1x 104 ~ 6x 104
3 05 (A) ~ 10° ~ 108
4 0, 1x 107? ~ 10°
5 NO, ~1x 107 ~ 1046
6 N0 ~ 6x 10° ~ 10°
295
7 HO, ~5x 10° ~ 23
8 OH ~ 5x 10° ~3
~~ 108
9 cH30, 10 100

39