United States
             Environmental Protection
             Agency
            Environmental Sciences Research
            Laboratory
            Research Triangle Park NC 27711
EPA-600 3-79-033
April 1979
             Research and Development
&EPA
A Problem with
Flux  Chamber
Measurements  of
Biogenic  Sulfur
Emissions

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                RESEARCH REPORTING SERIES

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                                            EPA-600/3-79-033
                                            April 1979
  A PROBLEM WITH FLUX CHAMBER MEASUREMENTS
        OF BIOGENIC SULFUR EMISSIONS
                     by

              Dian R. Hitchcock
            Hitchcock Associates
       Farmington, Connecticut  06032
           Contract DA-6-99-6358A
               Project Officer

               Lester Spiller
           Aerosol Research Branch
 Environmental Sciences Research Laboratory
Research Triangle Park, North Carolin  27711
 ENVIRONMENTAL SCIENCES RESEARCH LABORATORY
     OFFICE OF RESEARCH AND DEVELOPMENT
    U.S. ENVIRONMENTAL PROTECTION AGENCY
RESEARCH TRIANGLE PARK, NORTH CAROLINA  27711

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                                   DISCLAIMER

     This report has been reviewed by the Environmental Sciences Research
Laboratory, U.S. Environmental Protection Agency, and approved for publication.
Approval does not signify that the contents necessarily reflect the views and
policies of the U.S. Environmental Protection Agency, nor does mention of trade
names or commercial products constitute endorsement or recommendation for use.
                                       11

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                                    ABSTRACT

     Problems associated with identifying and quantifying factors that in-
fluence liquid-phase controlled evolution of hydrogen sulfide and organic
sulfides through the air-water interface are briefly reviewed.  It was found,
that at present flux chamber measurements of the release of these biogenic
substances from natural systems cannot be regarded as reliable estimates of
releases occurring when the system is not enclosed by a chamber.

     This report was submitted in fulfillment of Contract No. DA-6-99-6358A by
Hitchcock Associates under the sponsorship of the U.S. Environmental Protection
Agency.  This report covers the period July 2, 1976, to October 30,  1976.
                                      111

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                                  CONTENTS
Abstract
Figures
Tables .
     1,
     2,
     3,
     4.
Introduction
Methodology
Experimental
Conclusions
iii
 vi
 vi

  1
  2
  6
 17
References
                                                          18

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                                   FIGURES

Number                                                                Page

  1    Enhancement coefficient a for H^S transfer between seawater
       and air.  (The dependence of a on pH and T is similar for
       fteshwater transfer but calculated values are slightly lower
       than those plotted here*) .......... 	 . * .    9
                                   TABLES

  1    Freshwater H-S Exchange Constants (cm h~ )	   10

  2    Seawater H«S Exchange Constants (cm h~ )	   11
                                     vi

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                                    SECTION 1
                                  INTRODUCTION

     Biogenic sulfur gases may contribute significantly to the atmospheric load
of sulfur species.  This possibility has focused interest on the rate of emis-
sion of gases such as hydrogen sulfide  (H2S) and dimethyl sulfide  (CH3SCH3) to
the atmosphere from surface waters overlying sediments where they are naturally
produced, or from exposed wet sediments.  Several scientists have attempted to
measure these emissions using emission flux chambers, a method by which gases
released to the head space of a chamber placed over an emitting surface are col-
lected by air or some other suitable carrier gas passing through; the average
flux is then estimated from the measured concentrations of the gases of interest
in the carrier gas.  Such methods have been used by a number of researchers, in-
cluding Brannon (1973), Aneja (1975), Hill et_ al^. (1978), Adams e_t a±.  (1978),
and Hansen et^ al.   (1978) .

     The- purpose of this report is to call attention to the possible influence
of surface wind speed-induced liquid-phase turbulence on the evolution of^H^S
and other gases from water surfaces.  This study indicates that chambers which
isolate emitting water surfaces from ambient wind movements may inhibit the
evolution of H2S, CHsSCHs, and other biogenic sulfur gases from surface water.

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                                    SECTION 2
                                   METHODOLOGY

     The general problem with the flux chamber technique is that isolating the
emitting water surface will modify the exchange characteristics in ways pre-
sently impossible to fully evaluate.  Liss  (1973, 1975) explained some of these
problems in his discussions of the applicability of Danckw'aert's two-film model
of gas exchange through the air-water interface to gas exchanges between natural
waters and the atmosphere.  (See also Deacon, 1977, for a more general treat-
ment of these phenomena.)

TWO-FILM MODEL OF AIR-WATER EXCHANGE

     The'flux F of a gas through the air-water interface is the product of an
exchange constant K for the gas and Ac, the difference between the concentration
of the gas in the liquid near the interface and in the air above it:

                                    F = K Ac                        (Eq. 1)

     The model used by Liss provides that exchange takes place via molecular
diffusion through two stagnant films at the air-water interface:  a liquid film
and a gas film.  The overall exchange is described in terms of two transfer
coefficients:  the gas-phase transfer coefficient kg, and the liquid-phase trans-
fer coefficient k , as well as the gas gradients within each film layer.  How-
ever, as transfer between both films is limited by the slower of the two trans-
fers through each, the exchange is spoken of as being  'controlled' by the slow-
er of the two.

     If exchange is controlled by liquid-phase resistance  (assuming the gas
obeys Henry's Law), then the overall exchange constant will be K , where

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                                  KI   K±   Hk                      (Eq. 2)

and H is the value of Henry's Law constant at the temperature  (T) of interest.
If exchange is controlled by gas-phase resistance, then the overall exchange
constant will be K  , where
                  9
                                 1_ _ 1_   H_                       (Eq. 3)
                                 K    k    k,
                                  g    g    1
To predict the value of K, the overall exchange constant for a given gas from
these equations requires knowledge of their individual gas- and liquid-phase
transfer coefficients k  and k  for the circumstances in question.  Liss
(1973) has shown that the transfer characteristics of gases in wind tunnel ex-
periments may be predicted in terms of the gas-phase transfer coefficient
k       for water vapor, and the liquid-phase transfer coefficient k  .  .  for
oxygen:
                                kg(X) = rkg(H20)                    (E^ 4)
                                ki(x) = aki(o2)                     (Eq- 5)

where X is the species of interest, r is the ratio of the square roots of the
molecular weights of water and X, and a is an enhancement coefficient.  If X is
relatively unreactive and does not undergo rapid chemical reactions in water,
then a = 1.  However, if X is subject to very rapid chemical reactions, then
these will tend to increase its gradient in the .liquid film, enhancing its
transfer relative to 02.  The influence of enhancement will be a function of
the speed of the reaction and the thickness of the liquid film.  The most
important chemical reactions leading to enhancement of transfer through the
liquid film are dissociation and hydration.  For gases such as sulfur dioxide
(862) and H2S, for which these reactions are extremely rapid,

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                                         T-l

where T is the ratio of total to ionic forms of the gas in water at the pH of
interest.  For H2S in natural waters,

                                  =  [H2S] +  [HS~]                   (Eq. 7)
                                         [HS-]
because at pH ranges characteristic of natural waters, the sulfur is negligible.
     For gases such as water vapor and SC>2 , whose transfer is gas-phase con-
trolled, the value of k  varies with turbulence in the atmosphere near the
                       g
surface of the water.  Wind tunnel experiments have shown that k       varies
                                                                g
directly with surface wind speed u; the limited data available from field ex-
periments confirm these observations  (Liss, 1973, 1975, Hicks and Liss, 1976).
However, in the field, the influence of fetch on the development of turbulence
in the air near the air-water interface is important, and should be considered
when predictions of k  .   are made  (Hicks and Liss, 1976).
                     9 (X)

     The situation is considerably more complex for gases such as 02, carbon
dioxide  (CO2) , H2S, and most organic sulfides, whose transfer would seem to be
liquid-phase controlled in most circumstances.  Wind tunnel experiments indicate
that k. .  .  varies with the square of wind speed  (Liss, 1973, 1975, Deacon, 1977),
      1 ((J2>
but the paucity of field data makes it difficult to evaluate the applicability
of this model to the field.  One problem is that liquid-phase controlled transfer
is a function of turbulence in the liquid phase, and variations in surface wind
speed are only one mechanism influencing liquid-phase turbulence.  Turbulence
induced by currents, surface ripples, or other motions in the water may also
influence k.. .  . , as may differences between the temperature of the surface
           1(02)
microlayer and the layers immediately beneath it  (Liss, 1975, Deacon, 1977).
It is also evident that water turbulence may have different effects on  the ab-
sorption of a gas by the surface than on evolution  from.it.  Thus, Liss  (1973)
found that breaking the surface of the water from below influenced release of
02 from it more than absorption of 02 by it.  It is therefore evident that the

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perturbations in gas-phase transfer caused by a chamber will vary with the
manner in which it is used.  Furthermore, the limited information available
about these phenomena only concerns transfer between water and air; no in-
formation is available about transfer between the surfaces of wet sediments and
air.

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                                    SECTION 3
                                  EXPERIMENTAL

CALCULATING THEORETICAL LIQUID-PHASE EXCHANGE CONSTANTS

     The preceding discussion has shown that numerous uncertainties make it
difficult to fully evaluate the possible chamber-induced perturbations in
liquid-phase-controlled release of biogenic sulfides emitted from water or
other wet surfaces to the air.  However, some insight into the possible in-
hibition that may result from the use of flux chambers may be derived by com-
paring theoretical values of K  calculated from Equations 2 through 7 above as
a function of a hypothetical 'wind speed' inside and outside the chamber, using
for this purpose the wind tunnel results of Liss (1973).  There are, however,
many problems associated with this approach.

     Liss (1973) has determined the values of k.,. . , k. ,nn . , and k  .„ n. in
wind tunnel experiments as a function of wind speed vi for the range 1..6 to 8.1
m s   (where u is measured at 10 cm above the surface).  In this range, k
             —
and k .     are remarkably well described by the following linear equations:
     g(H2U>
                             k. ,. .  = 0.52 + 0.165 u2               (Eq. 8)
                                                   —
                            kg(H20) = 18'6 + 1136 H                 
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films determines the transfer characteristics.  It seems likely that in very
still water, other factors such as the composition and thickness of the sur-
face organic film microlayer may be more important (Liss, 1975).  I know of
no wind tunnel or field experiments relating gas emission rates to extremely
low wind speeds.  A second problem is that H2S and perhaps other biogenic sul-
fides appear to be released in quantity from wet sediments exposed to the
atmosphere.  Film thicknesses under these circumstances are unlikely to vary
with high speed in the manner implied by Liss's experimental data.  A third
problem is that there is always some gas-phase turbulence in the chamber, be-
cause the carrier gas is usually passed through the chamber in order to sample
its composition.  Some flux chambers employ a stirring device to ensure good
mixing of the emitted gases in the carrier gas  (Hill et^ al_. , 1978, Aneja et al. ,
1978) but others do not  (Brannon, 1973, Adams e_t al^. , 1978, Hansen et^ al_. , 1978).
Those who have used a stirrer have not reported the effects of stirring on emis-
sions.  Lacking any basis for an alternative procedure, I have therefore assumed
for the purposes of this illustration that 'wind speed1 inside the chamber is
equal to zero, and that Equations 8 and 9 may be extrapolated to u_ = 0.

     Some of the problems introduced by these assumptions, and the difficulty
of evaluating the effect of the use of chambers on H2S emissions are due to
the very large range of a values which are likely to be encountered in nature.
T and therefore a vary exponentially with pk  , the first apparent dissociation
constant for H2S, since

                               Pkl + Log'"[lfi]" = PH                 (Eq> 10)

and pk  varies with T and chlorinity in seawater, and with T and ionic strength
in freshwater.  Goldhaber and Kaplan  (1975) have shown that the value of pk
for seawater may be very accurately calculated  from the following equation:

                 pk, ,          = 2.572 + 1359.96/T - 0.169 C11/3    (Eq. 11)
                   1(seawater)
where T  is  °K and Cl  is  chlorinity  (°/00).   Their  work also  suggests  that
this equation may be  adapted  to  determine  approximate  values of pk for
freshwater  as a  function of ionic strength and  T by  assuming that freshwater
                                       7

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constitutes a very dilute solution of seawater.  If so, then Equation 12 will
be appropriate to many of the ionic strengths encountered in freshwater  (about
0.05):
                       Pkn if   u  4.  x = 2-35 + 1359.96/T           (Eq. 12)
                         1 (freshwater)

Because of exponential dependence of a on pH and T  (see Figure 1), the value of
k       may become very large at some pH and T values common in natural habitats.
 1(H2S)
Consequently, at theoretical 'zero1 wind speed when k  .   .  = 18.6, the
                                                     g(H2U)
theoretical value of K  .     is determined by that of k  .    .  The extrapola-
tions used in the above calculations imply that H2S transfer may become gas-
phase controlled inside the chamber, since
                                                                    (Eq. 13)
                          K1(H2S)     akl(02)   Hrkg(H20)
and the values of H employed for these calculations range from 0.2942 at 10°C
to 0.5462 at 35°C.  If this" procedure is correct, it implies that in natural
habitats, where gas-phase turbulence above the emitting surface may be low, H2S
transfer may be gas-phase limited.  Therefore, the question of the influence
of gas-phase turbulence on H2S transfer cannot be ignored in considering the
release of H2S to the.atmosphere, either in chambers or in the field.
     Tables 1 and 2 list some theoretical values of K       calculated from the
                                                     x(H2S)
equations listed above.  Values of K           are also calculated as a function
                                    i(CH3SCH3 J
of wind speed, assuming a Henry's Law constant of 0.3, as given by Liss and
Slater (1974).  (changes in the value of H with T will not influence the theo-
retical values of K  for this gas, because its transfer is always liquid-phase
controlled.)
     Examination of these tables confirms that the emission rates of neither
organic sulfides nor of H2S observed inside flux chambers should be assumed
to be representative of the emissions that might be occurring outside them.
The liquid-phase controlled exchange constant varies exponentially with wind

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                500 -
Figure 1.  Enhancement coefficient a for H2S transfer between seawater and
air.  (The dependence of a on pH and T is similar for freshwater transfer
but calculated values are slightly lower than those plotted here.)
speed.  In the case of unreactive organic sulfides, the apparent inhibition
(calculated as the ratio of the exchange constant at a selected speed to that
at 'zero' wind speed) ranges from about 3 at 2 m s"1 to about 16 at 6 m s"1.
In freshwater at pH values corresponding to relatively low concentrations of
hydrogen sulfide ions  (HS~) relative to aqueous H2S, the inhibition is similar
to that for an unreactive organic sulfide.  However, in freshwater at pH values
above about 8, and in seawater at all pH values tabulated, emissions inside the
chamber appear to be gas-phase controlled, and the  inhibition effects are con-
sequently larger and more variable, being critically dependent on a.  Further-
more, where a is an important variable, the effects of pH are large.  Thus, for
freshwater  (Table 1) at pH 8.0, the inhibition factors range from about 5 to 25
for u_ = 2 to 6 m s-1 at all temperatures shown.  At pH 9.0, these inhibition
factors range from 20  to about 100.  In seawater  (Table 2) the influence of pH
is similar, but the inhibition factors tend to be  larger.  Thus, at pH 7.8, the
inhibition  factors range from about 5 to 26, but at pH 9.0 they range from about
35 to 150.

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              TABLE 1.  FRESHWATER H2S EXCHANGE CONSTANTS (cm h"1)

u

0
2
4
6

0
2
4
6

0
2
4
6

0
2
4
6
K(CH3SCH3)

0
1
3
6

0
1
3
6

0
1
3
6

0
1
3
6

.4
.2
.2
.4

.4
.2
.2
.4

.4
.2
.2
.4

.4
.2
.2
.4

pH 6.
T = 10°C
0.
1.
3.
7.
T = 20°C
0.
1.
4.
8.
T = 25°C
0.
1.
4.
8.
T = 35°C
0.
1.
4.
10.

5
Pk
6
4
8
9
Pk
6
6
2
5
Pk
6
6
4
9
Pk
7
8
9
0

7
= 7
0
2
5
10
= 6
0
2
6
13
= 6
1
2
7
14
= 6
1
3
8
17

.0
.15
.7
.0
.4
.9
.99
.9
.4
.4
.0
.91
.0
.6
.0
.3
.76
.2
.2
.6
.5
K(H2S)
7.5

1.2
4
10
20

1.6
5
13
27

1.8
5.7
15.3
31.1

2.3
8
20
41

8.0

2.0*
9
25
50

2.8*
13
35
70

3.2*
15
41
83

4.2*
21
56
113

8.5

3.0*
26
68
136

4.0*
37
98
194

4.6*
44
116
230

5.9*
61
160
317

9.

3.
72
183
351

4.
103
261
496

5.
122
308
585

6.
167
421
794

0

6*




8*




4*




8*




*Transfer is gas-phase controlled.
                                      10

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               TABLE 2.   SEAWATER H2S EXCHANGE CONSTANTS (cm h~ )

K(CH3SCH3)
u

0
2
4
6

0
2
4
6

0
2
4
6

0
2
4
6

0
1
3
6

0
1
3
6

0
1
3
6

0
1
3
6

.4
.2
.2
.4

.4
.2
.2
.4

.4
.2
.2
.4

.4
.2
.2
.4

pH 7.8
T = 10°C pk
2.0*
11
29
58
T = 20°C pk
3.0*
15
40
81
T = 25°C pk
3.4*
18
48
96
T = 35°C pk
4.4
25
65
132
K
8.0
= 6.88
2.6
16
43
86
= 6.80
3.5*
23
61
123
= 6.64
4.0*
27
72
145
= 6.49
5.2
38
99
200
(H2S)
8.2

3.0*
25
65
129

4.0*
35
93
184

4.6*
42
110
218

5.8
58
152
301

8.4

3.3*
37
97
192

4.4*
54
139
274

5.0*
64
165
324

6.3
88
228
445

9.0

3.8*
119
294
541

5.0*
169
415
758

5.7*
199
487
887

7.1
270
656
1185

*Transfer is gas-phase controlled.
                                      11

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     These tables also show that the behavior of organic sulfides and H£S in-
side a chamber cannot be readily interpreted as reflecting differences in their
relative rates of release outside the chamber.  At all T and pH values listed
in both tables, the exchange constant for H2S is larger than that for CH3SCH3,
but the amount of the difference varies with pH and T.  In freshwater, at pH
values up to 7.5, the differences are small.  At higher pH values, the differ-
ences become more pronounced, and they also increase with T.  Thus, the H£S
transfer constant calculated for H£S exceeds that calculated for CH3SCH3 by a
factor of about 5 to 10  (depending on T) at pH 8.0, and by a factor of 6 to 17
at pH 9.0.  The influence of pH and T are similar for seawater, but the differ-
ences are larger.

     These calculations indicate a possible serious problem when interpreting
flux chamber observations.  As numerous uncertainties can be associated with
the behavior of gases released from surface water and/or wet sediments, these
results must not be viewed as constituting accurate descriptions of the proces-
ses occurring inside flux chambers, nor as providing useful predictions of the
exchange process in the natural habitat.

     These calculations do suggest, however, that in nature the release of H2S
from natural waters is an extremely complicated process.  If H2S release is as
sensitive to pH, T, and liquid-phase turbulence as these calculations imply,
then in some habitats actual releases must be very sensitive to short-term
variations in turbulence induced by wave action, currents, and perhaps even
disturbances of the surface due to biological activity.  The rate of supply of
H2S to water near the surface is probably not as variable as its release from
the surface to the air.  It is possible that the concentration of H2S in sur-
face water could be highly variable, building up to relatively high levels if
liquid-phase turbulence is minimal, and dropping rapidly as a response to emis-
sions resulting from short-term disturbances of the surface.  We may speculate
that disturbances associated with putting a chamber in place could deplete H2S
concentrations significantly, and would be followed by low measured emissions
until concentrations build up again.  If so, then the total duration of the
observations may be critical.
                                       12

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     These calculations also show that the relative concentrations of organic
 sulfides and H2S in surface water will rarely indicate the strength of their
 emissions.  Especially when pH and T are high, a low concentration of H2S may
 support a very large emission to the atmosphere, while much higher concentra-
 tions of CHsSCHs may yield much lower emissions.  These considerations throw
 doubt on the conclusions reported by Rasmussen  (1974), who inferred that CH3SCH3
 emissions from a small pond greatly exceeded H2S emissions, because the CH3SCH3
 concentrations were often one to two orders of magnitude higher than the H2S
 concentrations.  Examination of the original data (Bechard, 1974) shows that
 the values of the enhancement coefficients for H2S were usually in the range
 1000 to 3000, and that on the few occasions when the enhancement coefficients
were less than about 10, the H2S concentration exceeded the CH3SCH3 concentra-
tion.  Viewed in this light, the observations reported by Rasmussen imply that
the pond was a much larger source of atmospheric H2S than of CH3SCH3.

TEMPERATURE AND PH VARIATIONS IN NATURE

     A further cause for concern is associated with the influence of T and pH
on H2S evolution from surface waters, and the natural variation of these
parameters in habitats of interest.  Temperature and pH vary widely throughout
the diurnal cycle in most habitats where H2S is produced in abundance by the
metabolism of the bacterial sulfate reducers, and from which it may escape to
the atmosphere.  In natural aquatic systems, pH varies with the concentration
of C02, which in turn varies diurnally with the ratio of respiratory CC>2 pro-
duction to photosynthetic CO2 consumption.  The more productive the habitat,
 the greater the diurnal range of pH values.  In relatively shallow bays, creeks,
marshes and similar aquatic systems, T varies diurnally with surface heating,
 and reaches its highest values in synchrony with pH.  Skirrow  (1965) measured
 summertime diurnal T and pH ranges from 28°C and 8.0 at night  to 35°C and 8.9
 the next day in highly productive water in Redfish Bay  (Texas).  The correspond-
 ing values of a are 42 and  258.  In winter, the observed range was from 18°C
 and 8.0 at night to 23°C and 8.4 the next day  (a equal to 29 and 54).  In salt
 marsh pools and tidal creeks, the range of T and pH values may be greater, be-
 cause plant productivity may make more demands on the supply of CC>2, and the
 very shallow waters get hotter than  the surface waters of deeper bays.  However,

                                      13

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in these habitats the temporal cycles are liable to be more complex, due to
tidal effects.  Very large pH and T variations are also likely to occur in some
freshwater habitats.  Bechard (1974) reports 67 measurements of pH, T, and sul-
fur gas concentrations in a freshwater pond; T values greater than or equal to
25°C were observed in 20% of these, and pH values greater than or equal to 9.5
in 30%.

     Temperature is usually reported by many who employ the flux chamber tech-
nique, but not the pH or sediment or water H2S concentrations.  (Brannon, 1973,
and Hansen et_ al^. , 1978, are exceptions.)  As a result, one cannot determine
whether the flux measurements some researchers report reflect variations in the
relative productivity of H2S (and/or other biogenic sulfides), variations in
pH, or both.  Consequently, even if the chamber method were demonstrated to have
little significant effect on the transfer coefficients for these or other gases
in some habitats of interest, many results reported in the literature could not
be extrapolated to other times when the emitting surface may have other T and pH
values.

     These studies may be contrasted with those of Hansen et al.  (1978), who
used a flux chamber to study the emission of H£S from two tide pools.  They
determined H2S emissions at hourly intervals for durations of at least 24 hours,
and monitored light intensity, surface water pH, T, H2S and 02 concentrations,
and redox potential gradients in both water and underlying sediment as a func-
tion of depth.  In addition, they characterized the sediments in biologically
meaningful ways  (sediment grain size, organic matter content and distribution,
H2S and sulfate  (SOi*2 ) concentrations, and biological H2S production as a
function of sediment depth).  These authors reported the observed H2S emissions
in relevant units  (VIM m~2h~l, rather than g m~2y~ ) and related the variations
in emissions to variations in other relevant parameters.  The observed emissions
were accurately characterized as "minimum emissions from stagnant water".  In
consequence, this study provides much valuable information, even though the
HoS emissions may not be representative of emissions outside the chamber.
                                       14

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 POSSIBLE  EFFECT OF H2S EVOLUTION INHIBITION ON ITS  AQUATIC FATE
      Unlike  CHsSCHa  and most other organic  sulfide  gases  which are stable in
 water,  H2S may be oxidized by aqueous  02-   The  reactions  involved have not been
 studied in detail, but the limited evidence available  indicates that they are
 extremely  complex.   Reported H2S half-lives in  natural waters range from about
 20  minutes (Ostlund  and Alexander, 1963)  to over 100 hours (Skopintsev ejt al. ,
 1964),  but 2 to about 20 hours appear  to  be the most common values (Cline,  1968,
 Cline and  Richards,  1969,  Bella,  1975,  Wheatland,  1954, Almgren and Hagstrom,
 1974) .   Hydrogen sulfide aqueous oxidation  may  be  catalyzed by a number of
 transition series metals,  and inhibited or  catalyzed by organic substances
 (Chen,  1970, Cline,  1968,  Cline and Richards, 1969).  There is good evidence
 that  the uncatalyzed reaction occurring in  distilled water is less than first
 order in sulfide and greater than first order in 02 (Chen, 1970), and suggestive
 evidence that this may be the case in  some  natural  seawater systems (Broenkow,
 1969).   The  oxidation products include sulfite, sulfate,  sulfur, thiosulfate,
 and a number of polythionates; the relative abundance  of  these depends in part
 on  the  ratio of H2S  to 02 (Chen, 1970,  Cline, 1968, Cline and Richards, 1969).

      On the  basis of these reports one may  speculate that, if inhibition of
 H2S transfer to the  atmosphere occurs  in  flux chambers, then a greater pro-
 portion of the H2S present in the water may be  lost through aqueous oxidation
 than  would otherwise be the case.  It  is  possible that H2S oxidation products
 could then react with organic substances  present in sediments and surface
 waters  where H2S is  produced to form organic sulfide gases (Hitchcock, 1976) .
 The latter speculation is further strengthened  by the  results reported by Hill
 et_ al.  (1978) , and Adams et^ al.  (.1978) , who found that CHaSCHs was released in
                                                      — r\ ™ 1
 relatively large quantities (on the order of 0.1 mg m  h M  in the daytime from
 Spartina alterniflora sediments.  Rasmussen (1974)  reports  that he found high
 concentrations of CHaSCHs in water or in the head space of  cultures of decaying
 freshwater algae species when the use of a cotton stopper permitted some entry
.of 02 ,  but only H2S when the culture was isolated from the  atmosphere and the
 oxygen  concentrations in the head space dropped (see also Lovelock et al. ,
 1972).
                                       15

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     The authors of these reports attribute the CH3SCH3 to the bacterial decom-
position of organic sulfur present in decaying plant matter.  However, this ex-
planation seems unlikely on quantitative grounds.  Organic sulfur in plants is
present only in trace quantities  (chiefly in amino acids); during decomposition
very little of it is released as an organic sulfide gas.  Lovelock et al.  (1972)
measured CHsSCHs release rates in a number of substances.  These rates, correct-
ed for organic sulfur concentrations in the substrate, imply release rates of
about 10 i:ig g 1 of organic sulfur h l or less.  This, in turn, suggests that
the average CHsSCHs release from organic sulfur decomposition in Spartina sedi-
ments may be about 0.001 yg m 2h *, assuming organic sulfur concentrations in
                            *"* 0
soils are as high as 100 g m   (an estimate that assumes a very large organic
sulfur standing crop in bacterial tissue).  The production rates are far too
low to account for the observed organic sulfide releases; therefore, their
source must be sought for elsewhere.  The only other source is the H2S which
is produced in these soils in enormous quantities by bacterial sulfate reducers.

     The dependence of CH^SCt^ concentration or release on 02 availability,
observed by Rasmussen (1974) and Lovelock et^ a_l_. (1972), point to a possible
source:  '3 reaction between an H2S oxidation product and other constituents
normally present in muck soils and other habitats where bacteriogenic H2S is
produced in quantity.  Thus, sulfite, polythionates, or elemental sulfur may
react with organic acids or hydrocarbons to yield CH3SCH3 and possibly other
organic sulfides in the quantities necessary to support the observed emissions.

     This line of reasoning suggests that inhibition of H2S evolution from
surface waters or even from surface sediments could promote its aqueous oxida-
tion.  This process could conceivably result in the partial transformation of
H2S into organic sulfides.
                                      16

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                                    SECTION 4
                                   CONCLUSIONS

     The foregoing discussion has demonstrated that there are great uncertain-
ties regarding the use of flux chamber measurements to estimate the flux of
biogenic gases from aquatic habitats to the atmosphere.  These include the na-
ture and magnitude of the effect of chambers on the transfer constants describ-
ing release to the atmosphere of H2S and inert gases.  The extreme variability
of T and pH in natural habitats of interest also makes extrapolations from
chamber observations to other times extremely questionable.

     Until research on mechanisms influencing liquid-phase controlled transfer
of gases from water surfaces and wet sediments demonstrates otherwise, flux
chamber measurements cannot be regarded as providing reliable estimates of the
rate of evolution of H2S or other biogenic sulfides to the atmosphere.  For the
present, chambers should be regarded as useful ancillary tools to be used in
the context of well-controlled studies that characterize and carefully monitor
the physical and biological characteristics of the systems being examined.
Furthermore, the concentrations of gases in the sediment or water under study
should be monitored together with T, pH, and the very numerous other chemical
properties known to influence the biological production and chemical fate of
bacteriogenic sulfides.
                                      17

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                                   REFERENCES

Adams, D.F., S.O. Farwell, N.R. Bamesberger, and A.E. Sherrard.  1978.  Measure-
     ment of Biogenic Sulfur-Containing Gas Emissions from Soils and Vegetation.
     Paper presented before the Air Pollution Control Association Annual Meet-
     ing, Philadelphia, PN.  June.

Almgren, T., and I. Hagstrom.  1974.  The Oxidation Rate of Sulphide in Seawater.
     Water Res.  8:395-400.

Aneja, V.P.  1975.  Characterization of Atmospheric Sulfur Compounds.  Master's
     Thesis.  North Carolina State University, Raleigh, NC.  144 pp.

Aneja, V.P., J.H. Overton, Jr., and L.T. Cupitt, 1978.  Direct Measurements of
     Emission Rates of Some Atmospheric Biogenic Sulfur Compounds.  Presented
     at the 176th Meeting of the American Chemical Society, Miami, Fl.
     September 11-14.

Bechard, M.J.  1974.  Emission of Volatile Organic Sulfides by Freshwater Algae.
     Master's Thesis.  Washington State University, Pullman, WA.  53 pp.

Bella, D.A.  1975.  Tidal Flats in Estuarine Water Quality Analysis.
     EPA-660/3-75-025, U.S. Environmental Protection Agency, Corvallis, OR.
     186 pp.

Brannon, J.M.  1973.  Seasonal Variation of Nutrients and Physicochemical
     Properties in the Salt Marsh Soils of Barataria Bay, Louisiana.  Master's
     Thesis.  Louisiana State University, Baton Rouge, LA.  130 pp.
                                      18

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Broenkow, W.W.  1969.  The Distribution of Nonconservative Solutes Related to
     the Decomposition of Organic Material in Anoxic Marine Basis.  Ph.D.
     Dissertation.  University of Washington, Seattle, WA.  207 pp.

Chen, K.Y.  1970.  Oxidation of Aqueous Sulfide by Oxygen.  Ph.D. Dissertation.
     Harvard University, Cambridge, MA.  170 pp.

Cline, J.D.  1968.  Kinetics of the Sulfide-Oxygen Reaction in Seawater:
     An Investigation at Constant Temperature and Salinity.  Master's Thesis.
     University of Washington, Seattle, WA.  68  pp.

Cline, J.D., and F.A. Richards.  1969.  Oxygenation of Hydrogen Sulfide in
     Seawater at Constant Salinity, Temperature, and PH.  Environ. Sci.
     Technol.  3:838-843.

Deacon, E.L.  1977.  Gas Transfer To and Across an Air-Water Interface.  Tellus
     29:363-374.

Goldhaber, M.B., and I.R. Kaplan.  1975.  Apparent Dissociation Constants of
     Hydrogen Sulfide in Chloride Solutions.  Mar. Chem.  3:83-104.

Hansen, M.H., K. Ingvorsen, and B.B. J^rgensen.  1978.  Mechanisms of Hydrogen
     Sulfide Release from Coastal Marine Sediments to the Atmosphere.  Limnol.
     Oceanog.  23:68-76.

Hicks, R.B., and P.S. Liss.  1976.  Transfer of Sulfur Dixoide and Other
     Reactive Gases Across the Air-Sea Interface.  Tellus 28:348-353.

Hill, F.B., V.P. Aneja, and R.M. Felder.   1978.  A Technique for Measurement of
     Biogenic Sulfur Emission Fluxes.  J.  Environ. Sci. Health 3:199-225.

Hitchcock,  D.R.  1976.  Chemistry of Sulfur  in  the Natural Atmosphere.  Present-
     ed at  the Workshop on the Chemistry of  Atmospheric Sulfur,  Drexel Univer-
     sity,  Philadelphia, PN.  October  12-14.
                                      19

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Liss, P.S.  1973.  Processes of Gas Exchange Across an Air-Water Interface.
     Deep-Sea Res.  20:221-238.

Liss, P.S.  1975.  Chemistry of the Sea Surface Microlayer.  In:  Chemical
     Oceanography, Vol II., J.P. Riley and G. Skirrow, eds.  Academic Press, NY.
     pp. 193-243.

Liss, P.S., and P.G. Slater.  1974.  Flux of Gases Across the Air-Sea Interface.
     Nature 247:181-184.

Lovelock, J.E., R.J. Maggs, and R.A. Rasmussen.  1972.  Atmospheric Dimethyl
     Sulphide and the Natural Sulphur Cycle.  Nature 239:452-453.

Ostlundj H.G., and J. Alexander.  1963.  Oxidation Rate of Sulfide in Sea Water,
     a Preliminary Study.  J. Geophys. Res.  68:3995-3997*

Rasmussen> R.A.  1974.  Emission of Biogenic Hydrogen Sulfide.  Tellus 26:254-
     260»

Skirrow, G.  1965»  The Dissolved Gases—-Carbon Dixoide.  In:  Chemical
     Oceanography, Vol. 1> S.P. Riley and G. Skirrow, eds.  Academic Press, NY.
     pp. 227-322.

Skopintsev, B.A., A.v. Karpbv, and O;A. Vershihina.  l964i  Study of the
     Dynamics of Some Sulfur Compounds in the Black Sea Under Experimental
     Conditions.  Soviet Oceanography 4:55-72.

Wheatland, A.B.  1954.  Factors Affecting the Formation and Oxidation of
     Sulphides in a Polluted Estuary.  J. Hyg.  52:194-210.
                                      20

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TECHNICAL REPORT DATA
(Please read Instructions on the reverse before completing)
i.

4.
REPORT NO.
EPA-600/3-79-033
TITLE AND SUBTITLE
2.


A PROBLEM WITH FLUX CHAMBER MEASUREMENTS
OF BIOGENIC SULFUR EMISSIONS
7.
9.


12
AUTHOR(S)
Dian R. Hitchcock


PERFORMING ORGANIZATION NAME AND ADDRESS
Hitchcock Associates
Norton Lane
Farmington, CT 06023




. SPONSORING AGENCY NAME AND ADDRESS
Environmental Sciences Research Laboratory - RTF, NC
Office of Research and Development
U.S. Environmental Protection Agency
Research Triangle Park, North Carolina 27711
15
16
SUPPLEMENTARY NOTES

ABSTRACT
Problems associated with identifying

3. RECIPIENT'S ACCESSION>NO.

5. REPORT DATE
April 1979



6. PERFORMING ORGANIZATION CODE
8. PERFORMING ORGANIZATION REPORT NO.
10. PROGRAM ELEMENT NO.
1AA603A AA-145 (FY-79)
11. CONTRACT/GRANT NO.
DA-6-99-6358A
13. TYPE OF REPORT AND PERIOD COVERED
Final 7/76 to 10/76
14. SPONSORING AGENCY CODE
EPA/600/09




and quantifying factors that influence
liquid-phase controlled evolution of hydrogen sulfide and organic sulfides through
the air-water interface are briefly reviewed. It was found that at present flux
chamber measurements of the release of these biogenic substances from natural


17
a.
systems cannot be regarded
system is not enclosed by
as reliable estimates of releases occurring when the
a chamber.



KEY WORDS AND DOCUMENT ANALYSIS
DESCRIPTORS

*Air pollution Test chambers
^Hydrogen sulfide
^Organic sulfides
^Biological productivity
*Chemical analysis
*Flux density
13. DISTRIBUTION STATEMENT
RELEASE TO PUBLIC



b. IDENTIFIERS/OPEN ENDED TERMS

19. SECURITY CLASS (This Report)
UNCLASSIFIED
20. SECURITY CLASS (This page)
UNCLASSIFIED
c. COSATl Field/Group
13B
07B
07C
08A
07D
14B
21. NO. OF PAGES
26
22. PRICE

EPA Form 2220-1 (9-73)
                                                           21

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