-------
-14-
16% from the combustion of petroleum products, primarily residual
fuel oil. The remaining emissions resulted from refining operations
(M%) and non-ferrous smelting (^10%) , A study on anthropogenic SO
production by the Study of Critical Environmental Problems (SCOP)
(M.I.T., 1970) concluded that globally 93 x 10 kg SO were produced in
the year 1967-1968. Kellogg et al. (1972) point out that this estimate
may be low because the "emission factor" used for this global estimate
is probably applicable only in the United States and will be higher
in other nations which rely more heavily on fossil fuels with high
sulfur contents such as coal, Kellogg et al, believe an estimate of
9
about 100 x 10 kg SO,, per year would be much more reasonable; of this,
almost 94% is emitted in the Northern Hemisphere alone. Friend (1973)
9
has estimated anthropogenic emissions of SO- at 130 x 10 kg per year.
There are, of course, natural sources of sulfur dioxide as well
but the total amounts derived from these sources are extremely difficult
to quantify. It is believed that most natural S0? is released by
volcanoes. Kellogg et al. estimate that the quantity released by
volcanoes is about two orders of magnitude less than the amount they
9
estimated to be a result of man's activities (about 1.5 x 10" kg/yr
vs.-1-0 x 10 kg/yr). There is, as the authors point out, a good deal
of uncertainty in their estimate. Stoiber and Jepsen (1973) estimated
9
annual volcanic emissions of SCL to be 15 x 10 kg, an order of magnitude
greater than the estimate of Kellogg et al„
There are no other widely accepted natural S0~ sources, Kellogg
et al„ have pointed out that despite the fact that SO is so very
soluble in water, sea water might be a source rather than a sink if the
-------
right physicochemical conditions prevail. Without further conv.i nc «>i:.
evidence though on this subject, one should probably contend thai- : <.> ;
water would, in to to, provide only relatively minor amounts of SO,, 11>
the atmosphere, if any at all
Removal Mechanisms
Sulfur dioxide is very soluble in water; at 20°C and 1 atmo;.j i'• •>
and either oxidizes to sulfate or photochemically reacts with other
atmospheric contaminants. Therefore, sulfur dioxide is removed lrn;,i
atmosphere by various mechanisms involving water or other compound <;
The major identified sinks for this gas are: precipitation scavrm-;"
chemical conversion, and absorption by soil, water, rock, and plan!
The lifetime of sulfur dioxide in the atmosphere is estimated to i.i->r
between 20 minutes and 7 days (Nordo, 1973). Following is a discus-.'
of the various mechanisms mentioned.
Vegetation: A portion of the much needed sulfur used by plan!-.
in metabolic processes has been shown by Fned(1948) to be atln ru>! >••!
to the direct absorption of S09 from the atmosphere, especially in •:
where the soil is sulfur deficient. The ability of the plant to r' >,''
this sulfur effectively without damage is dependent upon the rat'- oi
absorption of S0? and the rate of production of sulffl-tes The [I'anl
can utilize this sulfur only if the production of sulfites does uni
exceed the capacity of the plant to oxidize the latter to sulfate1-
Studies by de Cormis [1968) suggested that the extent or S00
absorption is directly proportional to the atmospheric SO,, cone en', ^ru
-------
-16-
and is not influenced by the amount of sunlight. Hill £1971) investigated
the uptake rates of several gases by an alfalfa canopy. His results
confirmed those of de Cormis. Hill found that SO was absorbed with a
deposition velocity of 2.8 cm/sec, given a wind velocity of 1.8-2.2 m/sec,
as well as a number of other fixed variables. He did note that, in general,
those gases readily absorbed by the alfalfa were those with the greatest
solubility in water [see Table II and Figure 1].
Table II- SOLUBILITY IN WATER AND UPTAKE RATE OF POLLUTANTS
Pollutant
CO
NO
°3
N02
S02
Uptake Rate
in Alfalfa*
2 9
[mol/m *sec) x. 10
0
2.1
34.7
39.6
59.0
Solubility
at 20°C
g/100 g
0,00234
0.00625
0,052
decomposes
10,8
''Concentration of the gas in the chamber was 2 x 10 mol/m ,
Using SO,, concentrations at 10 stations covering 5f>50 square
kilometers near Sudbury, Canada, as measured by Dreisinger and McGovern
(1970), and data from Figure 1, Hill calculated a hypothetical rate of
SO removal by vegetation (assuming a continuous cover of alfalfa) in
this area. He indicates that the area is capable of removing 40 pg SO /
2 3
m «sec or 560 x 10 kg/d, His analyses make a number of assumptions
-------
-17-
O
234
Pollutant Concentration
(mol/m3)X I06
Figure 1, Uptake rates of different pollutants by an alfalfa
-------
-18-
and yield a nice ball-park answer, but more important than the actual
value is the indication it gives as to how important vegetation can be
in cleansing our environment.
Factors which influence pollutant uptake by plant canopies have
been discussed by Bennett and Hill (1973). In a more recent communication
Bennett, Hill, and Gates (1973) present a model simulating pollutant
transfer between leaves and the free air surrounding it. This model
is based on the rate of exchange via a series of external and internal
leaf mass transfer resistances„ The model indicates the importance of
gas solubility within the leaf. Thus, use of deposition velocities
and vapor phase concentration alone, as done by various investigators
(Owers and Powell, 1974, Shepherd, 1974; Chamberlain, 1960; and Spedding,
1969) will be inadequate for the prediction of uptake rates of a gas
by plants. In addition, variables which influence momentum and mass
transfer boundary layer thickness, such as wind velocity and gas
diffusjlvity, are shown to be important in determining the rate of
pollutant uptake by plants. One should bear in mind from known mass
transfer theories that the shape of the plant and leaf would influence
the rate of uptake as well. The effect of shape, however, has not
been tested or determined.
Estimates of the amount of SCL removed each year by vegetal
9
absorption vary greatly. Eriksson's (1963) cycle estimates 75 x 10 kg
9
SO -S are removed in this manner. Junge (1963) estimates 70 x 10 kg
S02-S, Robinson and Robbins (1968) estimate 26 x 109 kg SO -S, Kellogg
et al. (1972) estimate 15 x 10 kg S02-S, and Friend (1973) also
9
estimates that 15 x 10 kg SO -S are absorbed by vegetation each year.
-------
-19-
Th e designation "SCL-S" is a means of expressing the amount of SCL
t- i
present as its equivalent elemental sulfur weight.
Soi1: Recent studies have shown that soils are capable of absorbing
significant amounts of sulfur dioxide. In 1936, Vandecave/e and associates
(as cited in Bohn, 1972) noted that acid soils absorb only small
quantities of SCL, Terraglio and Manganelli (1966), in studies of two
soil types, found not only that SCL was more readily absorbed by soil
with a higher moisture content, but that the reaction also appeared
to be dependent upon the pH of the soil, more SO being absorbed in
the soil where the pH was greater. It is, of course, not that simple.
The degree of absorption is also dependent upon such factors as the
mineral and organic consent of the soil, soil structure, ion-exchange
capacity and porosity, Faller (1968), as cited in Seim (1970), found
that the oxidation (a word Faller uses as synonymous with absorption)
of SCL was greatest in soils with the highest base saturation, other
factors being equal. Faller and Herwig (1970) suggest that the amount
absorbed is related to the exhangeable alkali and alkaline earth cation
content, and what is absorbed is converted to H«SO. when the water
content of the soil is sufficient (-20%).
Experiments performed by Seim (1970) showed that soils can be
important media for absorbing SO = He found that the amount of sulfur
absorbed by the 0-2 cm depth segment of a particular dry silt (Fayette
soil, analysis in his text) exposed to air containing 3720 pg SO /m
2
for 24 hours was equivalent to 10.6 g/m , He does note that it is
questionable whether the soil can maintain this absorption rate for
-------
-20-
extended periods of time, but it might be possible if sulfate is continually
removed from the soil by leaching and/or crop uptake,
Smith et al .. (1973) showed that soils of varying pH, organic
carbon content, and sand, silt and clay content all readily absorbed
sulfur dioxide, In a chamber containing air with an initial SCL
concentration of 270,000 yg/m (100 ppm), all six soils reduced the
SO content in the air by 95% in less than or equal to one minute when
the soils were dry, and in less than one-half minute when the soils
were 50% saturated. Sterilization of the soils caused no significant
change in their S0? uptake ability. The removal mechanism is
therefore probably not biological. Like Seim, Smith et al. suggest
that the SO absorbed is oxidized to sulfate which may then be subject
to leaching and/or plant uptake. If the sulfate content in the soils is
thereby reduced, the soils may maintain their SO absorption capacity.
This hypothesis has not yst been investigated.
Deposition velocities have often been used to determine the
removal of S07 from the air above soils and vegetation. From the data
available in the literature, deposition velocities for soils appear to
be less than those for vegetation,
Seim obtained deposition velocities for S07 of approximately 0.2
cm/sec for all the soils he investigated. Chamberlain (1960) reported
values of 0,3 and 0.7 cm/sec for average data over Great Britain. For
combined soils and plants, Chamberlain used a value of 1.8 cm/sec.
Spedding (1969) calculated a maximum deposition velocity of 1.5 cm/sec
for plants. Owers and Powell (1974) calculated a mean deposition velocity
of 0,8 cm/sec over Great Britain assuming the countryside is all grassland.
-------
-21-
Seiro points out that "if these values are valid and the 0,2 cm/sec value
obtained for soils alone in this study should prove to be valid over a
wide range of climatic conditions, the conclusion which must be drawn is
that plants account for greater absorption of SO from the atmosphere
than soil".
Estimates on the amount of SO- that is absorbed by soils are
lacking and, for the most part, missing from most sulfur cycles that
have been compiled„ Unless this process has been taken into account
in estimating the total sulfur deposited by dry deposition on the land
surface then one must conclude that there is an obvious omission in
9
the cycles. Eriksson (1963} estimated that 75 x 10 kg/yr of dry
9
deposition sulfur goes through plants into the soil while 25 x 10 kg/yr
are directly absorbed by the soil. Abeles et al. (1971), based on
experiments they themselves ran, concluded that soils of the United
9
States are capable of removing 4 x 10 kg of S0? per year,
Rock: Living matter has a threshold limit of tolerance to pollu-
tants, below which no injury will occur, and, in fact, the substance might
actually benefit from the presence of that pollutant. Rock reacts with
pollutants such as SO , at all concentrations. The results of these
reactions may not be visible for quite some time, for the effects are
cumulative. The extent of breakdown of the rock is therefore less a
product of the momentary concentration of the pollutant than the uptake
per unit time on a unit area of the material (Luckat, 1973).
The effects of SO- on frescoes, monuments and other edifices
have been most pronounced over the last century, expecially in Europe
where high-sulfur coal and oil are used as heating fuels. The basic
-------
-22-
destructive reaction is that of sulfuric acid (SO- + 1/2 CL + hLO -»•
1LSO.) on the carbonate matrix of limestone and sandstone in the presence
of moisture. Spedding (1969b) has shown that as the relative humidity
in the air increases, the SO- uptake rate by oolitic limestone increases
significantly (Table III).
Table III. LPTAKE OF S02 BY OOLITIC LIMESTONE
Relative
Humidity
11
13
79
81
so2
concentration
Ug/m3
360
.280
LOO
370
Time of
exposure
20
40
48
10
Uptake
yg S02/cm2
of surface
0.069
0.061
0.24
0.28
Uptake
rate
5.0
2.2
7.2
40.3
The product of the; reaction between sulfuric acid and the carbonate
matrix is gypsum if sufficient evaporation occurs:
CaCO- + H SO, + H00 -*• CaSO/ 2H.O + C0_
3242 42 2
Because the calcium carbonate is slowly being replaced by gypsum the
rock would be subject to increased weathering rates due to:
(1) an enhanced chemical disintegration caused by the much
greater solubility of gypsum in water than calcium
carbonate. The newly formed mineral would be subject
-------
-23-
to dissolution in water with accompanied leaching out
of the rock, and
(2) an enhanced physical disintegration caused by the almost
two-fold volume expansion in the rock accompanying the
formation of gypsum.
Other properties such as the density and porosity of the rock
are also important and affect the amount of weathering to be expected
in a rock per unit uptake rate; a rough, porous, lime-cemented
sandstone would be expected to weather faster than a smooth, dense
limestone (Luckat, 1973).
To determine whether the absorption of S07 by sedimentary rocks
would constitute a significant sink for SO , a number of assumptions
were made:
(1) Knowing that approximately 30% of the total earth's surface
14 2
area of about 5 x 10 m was land (Holmes, 1965), and
assuming that perhaps 5% of the land surface had exposed
rock (F. E, Wickman, 1974), and of this approximately 75%
is sedimentary (Leet and Judson, 1965), it was calculated
that about 1% of the total earth's surface was covered
with rock capable of absorbing SO .
(2) Luckat (1973) and Spedding (1969b) found SO absorption
2
rates ranging from 5-200 mg/m -d in sandstone and limestone,
respectively. Luckat's measurements were taken in a highly
industrialized region of Germany; Spedding's measurements
were made with S0? concentrations approximately 100 times
greater than the average world-wide background concentration,
-------
-24-
2
For this calculation, the lower limit of 5 mg/m -d
was used,
9
From these values it was calculated that approximately 9 x 10 kg
n
SO /yr or 4.5 x 10"' kg S/>r could be removed by stone under these
optimal and exaggerated conditions. Comparison of this value with those
in Table V shows this value to be considerably smaller than that of any
other natural SO sink.
There are many problems related with making any such estimate as
was done above. First of all, any estimate of the total percentage
of the earth's surface which has exposed rock is entirely speculative
at this point in time because geological maps are not available for all
parts of the world. Then again, there is the problem of defining
an outcrop and mapping it in its strictest sense; that is, mapping the
outcrop without magnification and without the inclusion of soils or
detritus as part of the outcrop. There is also the fact that not all
the sedimentary rock that is exposed is sandstone or limestone; much
of it is shale or mudstone or something of similar density which would
probably not absorb SCL to any extent. Finally, the areas where most
of the rock does outcrop would be in areas where the S0_ level would
probably be quite low. It can therefore be assumed that rock constitutes
a negligible sink for SO on a global scale,
Water Bodies: Theoretical arguments in support of the contention
that sea water is capable of absorbing significant quantities of sulfur
dioxide from the atmosphere first appeared in a paper by Liss (1971)
and were modified by Liss and Slater (1974). Briefly, the rate at which
a gas is exchanged across an interface can be obtained by using a mass
-------
-25-
transfer coefficient, k, having velocity dimensions and defined as:
k = flux of gas
concentration difference l }
over a thickness, z
The overall mass transfer coefficient, K, based on liquid (£) or
gas (g) phase concentrations is related to individual mass transfer
coefficients by the following equation
1=1+1 I = 1 M.
K0 k0 + Hk °r K k + k0 ( J
£ £ g g g £
R£ = r£ + rg or Rg = ^ + rg (3)
where k,, and k are the mass transfer coefficients for the gas and
o
liquid phases respectively. The terms r,, and r are measures of the
o
resistances in the liquid and gas phases and are obtained from the
reciprocals of individual mass transfer coefficients. The Henry's law
constant, H, is defined as:
_ equilibrium concentration in gas phase (g/cm air)
equilibrium concentration of un-ionized dissolved gas in '
liquid phase (g/cm^ H-O)
and differs for every gas.
Whether the net transfer of a gas is controlled by the liquid
phase or gas phase is determined by the ratio of the resistances of
the gas to liquid phase (r /r,,) . For most gases (N_0, CM. and CC1 tu
name just a few) r /rc < 1, and transfer is controlled by the liquid pha-
g *
-------
-26-
Suifur dioxide transfer though is evidently controlled by the gas phase;
that is r /r0»l. For water vapor r. = 0 and the exchange across
g x, x,
the interface is controlled by processes in the gas phase. With the
knowledge of mass transfer coefficients for one contaminant, one can
approximate the mass transfer coefficients for other contaminants.
Liss and Slater (1974) used the mass transfer coefficient for
water vapor to calculate the overall mass transfer coefficients for
a number of gases crossing the air-sea interface. The results are
given in the following table::
Table IV. MASS TRANSFER COEFFICIENT FOR A NUMBER OF GASES
CROSSING THE AIR-SEA INTERFACE
k <
Gas g , H
cm/ sec cm/ sec
K*
g £ cm/sec
S02 0.45 9.6 3.8 x 10"2 573 0.45 (g)
NO 0.53 0.0055 1.6
CO 0.67 0.0055 50
CH, 0.885 0.0055 42
4
HO 0.833 «>
6.6 x 10 3 0.0055 (£)
-4
1.7 x 10 0.0055 (£)
-4
1,5 x 10 0.0055 (£)
0.83 (g)
*The overall exchange constant, K, is expressed on either a gas (g) or
liquid (£) phase basis.
Liss (1971) first showed that the exchange of S0? with aqueous
solutions was a function of pH (see Table V). A chemical enhancement
factor, a, equals the ratio of the kfl values for a reactive and an
A/
-------
-27-
Table V. CALCULATION OF THE RATIO OF THE GAS TO THE LIQUID
PHASE RESISTANCE
pH
2
2.8
3
4
5
6
7
8
9
a
2.7
11.7
18.0
169
1,376
2,884
2,966
2,967
2,967
k£(S02)
cm/sec
0.0075
0.0325
0,050
0.469
3,822
8.011
8,239
8.242
8.242
sec
130
30.
20.
2.
0.
0.
0.
0.
0.
cm
8
0
1
26
125
121
121
121
sec cm £
32 0, '
32 I,M
52 1 f.
32 15 1
32 123,1
32 25b, I)
32 264 L
32 26; ',
32 264,5
inert gas exchanging under identical conditions. As can be seen, abov<
a pH of 28 the gas phase resistance becomes dominant, and therefore
the exchange of SO with natural waters (pH 4-9) is controlled by 1 IK
gas phase. At a pH of 8 (approximately that of sea water) the gas pi' "
resistance has increased to more than 99 percent of the total resistm^
Brimblecombe and Spedding (1972) experimentally confirmed these resiiHr
of LisSc Liss and Slater (1974) suggest that the exchange of" othei
gases, such as NH , SO and HC1 which also are very soluble and
O »J
undergo rapid hydration reactions, might also be controlled by the
-------
-28-
resistance of the gas phase, It should be pointed out that the difference
in the values of the ratio r /rff for SCL at a pH of about 8 between
g )6 z
Tables IV and V is due to (1) a smaller chemical enhancement factor used
in calculations for Table IV yielding a slightly larger k0 ,„ , value,
2
and (2) a difference in calculating the value of k .„ ,. Whereas Liss
and Slater derived k , , by correcting the value of k , , by a factor
equal to the ratio of the square roots of the molecular weight of water
to the molecular weight of SO (as in Table IV), Liss (1971) had assumed
tnat K f,, 0^ -K fr^r* -» «
g(H20) g(S02l
Spedding (1972) collected samples of sea water and, by measuring
the amount of SO absorbed by the water from an atmosphere containing
SO , determined the deposition velocity of S0? over sea water. He found
••"hat the deposition velocity of SO increased linearly as the gas flow
rate increased. The natural buffering capacity of sea water prevented
any drop in the pH f-8), and thus would not limit the amount of SO
absorbed. Earlier, Terraglio and Manganelli (1967) had found a rapid
and substantial decrease in pH when SO was absorbed by distilled water.
Most data on SO solubility in the literature have been determined
when atmospheric S0? concentrations in the experimental chambers far
exceeded those found in ambient air. Hales and Sutter (1973) worked
towards closing this obvious gap by running a number of experiments
to help "quantify the relationships between S09 solubility, concentration,
and hydrogen-ion impurity at levels normally encountered in nature."
The dissolution of S0~ was assumed to proceed according to the reactions
set forth by Falk and Giguere (1958) :
-------
-29-
S0? + 2H20 + H30 + HSO~ (6)
"aq
HSO~ + H20 J H30+ + S03~ (7)
The second ionization (reaction 7) was assumed to be negligible.
Based upon the first two of the above reactions, Hales and Sutter
derived an "extrapolation equation, relating the concentration of total
dissolved SO- in water (c n ) to airborne concentration and solution
2. oU«
acidity" as follows:
[SO.] -[H,0+] + /[HJD+I2 + 4 K. [S00] /H
1 2J_g_ L 3 Jex L 5 Jex 1 L 2Jg
H
where those terms bracketed are concentrations in moles per liter and
[H 0 1 is the "excess" hydrogen ion concentration, defined as the
O C.A.
concentration of hydrogen ion in solution present due to sources other
than the dissolving of SO . K is the equilibrium constant for reaction ((>)
and H is the Henry's law constant. Extrapolation of SCL solubility data
from Johnstone and Leppla (1934) down to low ambient SO concentrations
show deviations ranging from 0.7 to 21.2% of those by Hales and Sutter,
In general, the percent deviation increased as the concentration of SO
in the gas phase decreased. The authors suggest that even though this
deviation does exist, the ability of equation (8) to predict low SO-
concentration solubility appears excellent and they recommend its use
whenever low concentration solubility data are needed.
-------
-30-
A more exact equation for determining the total dissolved SCL in
water can be derived based on the solution equilibria involved. The
resulting equation not only takes into consideration the pH of the solution
and the atmospheric partial pressure of the gas, P , but applies to
oU--i
different aqueous phase conditions, including both sea water and rain
water. Also it eliminates the use of the often confusing and now outdated
hydronium ion (H,0 ) . The dissolution and dissociation reactions for S09
-------
-31-
where [SO . H,,0] is the activity of solvated sulfur dioxide and Pqn is
the SO partial pressure in atmospheres. The defining equations for
the ionization constants are:
[H] [HSO~]
K =
l [S02-H20]
[H+] [SOI]
K = - — (14)
[HS03]
Using equations (12), (13) and (14), a molal mass balance for the system
can be written
aq
wherein, „ p
S02
mSO = ~Y -
bU2 YS09
aq 2
n aq
Ki H Pso2
•HSO-, - 7— rr-T,
* YHS03 L J
K2 Kl
m =
3
Substituting the values for the molality for each of these species
into equation (15), the following general formula for the concentration of
dissolved S09 is derived:
-------
-32-
2 2 2aq 3
The definition of pH is
pH = -log [H+]
Therefore, by convention, in equations the activity of the hydrogen ion
+ -r>H
can be written [H ] = 10 *' as desired.
For use in predicting SO,, solubility, equation (19) requires values
for H, K , K and the three activity coefficients in addition to the
pH and the SO partial pressure. Scott and Hobbs (1967) give H = 1.24,
K = 0 0127, and K0 = 6.24 x 10~ at 25°C, Corresponding data at other
temperatures are given by Johnstone and Leppla (1934). One case of
considerable interest is the absorption of S0~ in sea water. Activity
coefficients of aqueous sulfite species, SO «H?0, HSO_, and S0~ in sea
water are not available. However, Reardon (1974) determined activity
coefficients for the analogous carbonate species and also sulfate ion in
sea water. From a point of view of similarity of species and coherence among
the variables, the sulfate activity coefficient is rejected in favor of
use of the carbonate species activity coefficients. Hence,
YS02 S YH2CO°
aq
-------
•O-S-
By substituting these values into equation (19), and assuming that sea
water has a pH of 8.1, equation (13) reduces to
Lmcn = 8,42 x 107 Pcn (20)
b(J 2
where the amount of dissolved SO is seen to be a linear function of the
atmospheric partial pressure of SO-.
The minimum background concentration of S0_ is 1 yg/m which
corresponds to 3.8 x 10 atmospheres SO partial pressure. From
equation (20) the amount of S09 dissolved in sea water at equilibrium
_2
would be 3,2 x 10 mol/JL As an upper limit one might consider ambient
SO- concentrations founa by Luckat (1973) in highly industrialized
sections of Germany. The observed 360 yg/m is equivalent to P =
-7 2
1,4 x 10 atm. In equilibrium with such an atmosphere sea water
would absorb 11.6 mol SO /£. This molarity is an order of magnitude higher
than that where Henry's law is known to hold. Thus, the predictive
equation might fail under these circumstances.
When the dissolution of S02 takes place in rain water, the activity
coefficients required for equation (19) may be assumed equal to unity
because the ionic strength of rain water (a measure of the interionic
effect resulting primarily from electrical attraction and repulsions
between the various ions) would be very low, probably on the order of
-4
10 . Therefore, equation (19) reduces to
9 — 1 0
' "
[H ] [H ]2
-------
-34-
ni this case the hydrogen ion activity is dependent upon the amount
of SO absorbed and cannot be specified a priori. The charge balance
for the dissolution and dissociation reactions of S02 in water is
= [HSO~] + 2[SO] + [OH"]
By substituting in the respective expressions for the [HSO ] and [S0~]
o o
as given by equations (12, 13, and 14), and setting [OH~] = 10 /[H+],
an expression is derived for the partial pressure of S0 as a function of
+
[H] as follows:
[H+]3 - 10~14
p = ---
_
-? + -q
2 1.58 x 10 [H ] + 1,97 x 10
Choosing an initial value for the [H ] , and substituting that value
into equation (22), a corresponding value for Pcn (atm) can be found.
2
These values for Pcn and [H ] are then in turn substituted into
oUrt
equation (21) to find the amount of total dissolved SCL,
In pure rain water the pH is 7.0,. Any SO,, which dissolves would
produce an acid solution and a drop in pH. Thus, for pH = 7 , P =
2 -9
Zmcn = 0,. For a pH of 5, PQO and Sm are calculated to be 6.3 x 10
S02 bU2 bU2
atm and 1.0 x 10 mol/£, respectively,, For a pH of 3, P = 6.35 x
ou«
-5 -5
10 atm and £mcn = 1,1 x 10 mol/£. Clearly as Pcn increases, the
oU, ou_
L. <-
amount of dissolved SO increases and the pH decreases. The dissolved
S0? predicted by equations (21) and (22) agree excellently with the
experimental data of Hales and Sutter (1973) and Terraglio and
Manganelli (1967),
-------
Estimates of the amount of SO absorbed annually by the oceans
indicate the importance of this sink, Liss and Slater (1974) estimate
this SCL flux at 1,5 x 10 kg/yr based upon their own calculations.
Their estimate is in good agreement with those of Eriksson (1963) (2 x
TO11 kg/yr) and Robinson and Robbins (1968) (0.5 x 1011 kg/yr), but is
lower than that determined by Spedding (1972) (9.6 x 10 kg/yr). Liss
3
and Slater explain that this discrepancy is due to a 3 yg/m difference
in the mean atmospheric SCL concentration used by each and because
Spedding's value for the total resistance of the gas phase (1/K ) was
to
much lower. As noted earlier, Kellogg et al. (1972) consider the net
flux of S00 from air to sea to be negligible, based on observations made
by Pate et al. (person^ communication to Kellogg et al.) that in some
areas, where the equilibrium vapor pressure of SCL in surface waters
exceeds the partial pressure of SCL in the air above it, the ocean
might actually be a source of SO . Although this condition may exist
locally, on a global basis it would make far more sense to consider
the ocean as a sink in light of the high solubility of SCL in water.
Also, whereas laboratory experiments have shown that the solution and
rate of oxidation of S0_ to sulfate in distilled water is limited by
the pH as the solution becomes more acidic (Terraglio and Manganelli, 1967),
this would not be the case in the ocean (Spedding, 1972). The ocean has
a natural buffer capacity that maintains the pH at about 8.1- Therefore,
pH would not be a limiting condition in sea water and so it should have
the capacity to constantly absorb S02- This seems consistent with the
discovery of Liss (1971) and Liss and Slater (1974) that it is the gas
phase and not the liquid phase that limits the transfer of S0? across
the air-sea interface.
-------
-36-
Washout and Rainour, of SO : The major portion of SO present in the
atmosphere is probably removed by the processes known as rainout and
washout. Rainout invokes the scavenging of SCL and sulfate particles within
the clouds while washout, is the removal of these sulfur compounds below
cloud level via precipitation. The SO scavenged will undergo a series
of reactions, some catalytic, and ultimately form H~SO. drops or a
sulfate salt. The following paragraphs, due to their nature, might just
as well be included in the section on chemical reactions involving S0?
in the atmosphere, but they are included here so the reader can best
follow the sequence of events from the time the gaseous SO reacts with
water vapor in the atmosphere until the sulfate salts are removed by
precipitation or dry deposition.
From the time SO,., is absorbed by cloud droplets it is both innized
and oxidized by reactions (9), (10) and "(11) shown in the previous
section and (23)and (24) as follows (Miller and de Pena, 1972):
OH = HSO~ + H20 (23)
S0~ + 1/2 02 + S0~ (24)
dissolved
It is obvious then that any SO absorbed by water droplets will not only
contribute to the sulfate content of rainwater but will also cause a
decrease in the pH of the rain. The pH of rain usually ranges from 5-6
but can drop down to about 4 in areas where large amounts of SO are
present in the atmosphere. Because the pH drops, reaction (2s) is
important only in the first few seconds. Perhaps it would be easier
to visualize if the sequence above is given as the net reaction:
-------
-37-
S0o + H00 + 1/2 0_ -»• SO. + 2H+ (25)
/ 2. Z , . -. , 4
dissolved
Because the oxidation of SO,, in the liquid phase does not occur at a
rate fast enough to account for the sulfate content found in rain,
investigations were undertaken to find an effective catalyzing agent.
Experiments have shown that of all tne metals found in the atmosphere,
Mn salts were the most effective in promoting SO oxidation (Junge and
Ryan, 1958; Johnstone and Coughanowr, 1958; and Matteson et al., 1969).
Although Mn salts are the best known catalyst for SCL oxidation, the concen-
tration of these salts in the atmosphere is still not great enough
to account for the sulfate content in rain. Recent investigations though
have shown that when ammonia is present the rate of sulfate production in
solution is greatly enhanced (van den Heuval and Mason, 1963; Scott and
Hobbs, 1967; and Miller and de Pena, 1972),
The results of Scott and Hobbs'initial study as to how the
concentrations of SCL and NH effect the amount of S0~ produced and the
pH of the precipitation are shown in Table VI, The results are given
in terms of the SO? concentration because the rate of SO. production
depends only on [S0~] (reaction 22) The first line of Table VI gives
the initial values of the pH and the S0~ concentration corresponding
to partial pressures of SO and NH typical of those found in the
Earth's atmosphere, namely 20 ug/m (7 x 10 atm) and 5 yg/m (7 x 10
atm), respectively. The following lines show the effect of varying the
concentrations of S02 and NH . One can see that more sulfite is produced
given the same SO,, concentration when NH is present. Also, as the
concentration of NH increases so does that of the sulfite produced.
-------
-38-
Table VI,, INITIAL VALUES OF THE pH AND THE CONCENTRATION OF SULFITE
(SO*) ION FOR VARIOUS PARTIAL PRESSURES OF SO AND Nil..
o 2. . t
Partial Partial
pressure pressure Initial
of S0? (atm) of NH (atm) pH
*•• O
1 x 10~9 '7 x 10~9 6.34
7 x 10~9 0 4.97
1.4 x 10"8 7 x 10~9 6.21
7 x 10"9 1,4 x 10~8 6»48
5 x 10"6 5. x 10"6 6,35
5 x 10"6 0 3.55
Init j al
concentration
of SO'
(moles liter )
3.4 x 10~5
6.0 x JO"8
3,, 6 x 10 "
6.2 x i(f J
2.4 x 10"2
6,2 x 10"8
One can also see the effect that the presence of NH has on the pH, When
o
no ammonia is present the pH is highly acidic. When present though, the
hydroxyl ion produced upon the dissolution of ammonia in water tends to
neutralize the hydrogen icns produced by the dissolution of SO , Therefore
Zj
when ammonia is present the pH remains a little over 6, These values
have all been attained with C0? present in the gas phase at an average
concentration of 362,000 yg/m (311 ppm-). Therefore the acid pH is a
result of dissolution of C0_ as well as SO?J but the variability is
mainly attributable to S02 and NH concentrations,
Field investigations by Beilke and Georgii (1968) indicated that
the absorption of gaseous S0? by rainout and washout accounted for '/5%
of the sulfate content in rain water and that scavenging of sulfate
-------
-39-
paiticles contributed only 25%. The model formulated by Miller an-1
contradicts those measurements of Beilke and Georgii, Miller and i-
show that the sulfate content of rainwater is much more dependent, on
scavenging of particles rather than SCL. In their model, the huli.'i
content of rain water near a highly concentrated SO plume of or. i ;••
moderate particle concentration, showed that the contribution of- ••<
the total sulfate concentration was 4 times less than that of ;>ui( •!'
particles. Miller and de Pena's model makes more sense, especJdM,
because all reactions involving S09 in the atmosphere ultimate!;- i
to the formation of SO. .
Appreciable effort has been devoted to the analysis of SO , <
ging by rain (Engelmann, 1968; and Fuquay, 1970). Field measureinu i
SO washout (Hales, Thorpj and Wolf, 1971) have demonstrated that Un
a significant accumulation of SO in the water drops, and consequs, >:i
calculations of the washout based on deposition velocities are jna-d
A comprehensive analysis of reversible washout based on the
interaction of raindrops with atmospheric contaminants has been PI-,
by Hales (1972), This analysis indicates the use of overall mass <••
coefficients for determining the washout. From this analysis, li j
apparent that the degree of success in determining washout rat^s •! •
on estimating mass transfer coefficients and solubility data, J1;;J«'
used mass transfer coefficients for the limiting cases of gas phu'j*
control or liquid phase control without internal circulation wJLJu1
the drop to obtain the upper and lower limits of gas washout raito.
Further study (Hales, Dana, and Wolf, 1973; and Hales, Wolf, and f);m •
1973) has led to the development of a mathematical model for jpr
-------
-40-
ground level concentrations in the rain as a function of location beneath
a plume under stable meteorological conditions.
Atmospheric Reactions Involving SO : Reactions involving S09 in
the dry state, not unlike those discussed above for SCL in the wet state,
are very complex. The most important reaction involved here is the
photochemical oxidation of SCL which takes place in polluted atmospheres.
Early measurements of the rate of photo-oxidation of SO made by
Gerhard and Johnstone (1955) are the most widely quoted. They found the
rate of SCL oxidation tc proceed from 0.1-0.2%/hour, Recently, Cox and
Penkett (1970) measured this oxidation rate in experiments using purified
ambient air and found rates ranging from 0.04-0,65%/hour. Analysis of
the gaseous composition in their test chamber found small amounts of
-4 3
hydrocarbon compounds (.04 -CL2x 10 mol/m , or, 0.1-0.5 ppm) &nd
nitrogen oxides (<10 yg/m , or, <0.005 ppm), and they suggest that their
presence might be responsible for the high observed photo-oxidation rates.
Renzetti and Doyle (1960) had earlier suggested that the rate of photo-
chemical aerosol formation (the end result of the photo-oxidation of SCL)
is greatly accelerated in the presence of olefinic hydrocarbons and
nitric oxide. Endow, Doyle and Jones (1963) and Harkins and Nicksic
(1965) have shown that the resulting aerosols consist almost entirely
of sulfuric acid droplets when the relative humidity is greater than
or equal to 50%.
Cox and Penkett (1971a) have given experimental evidence to support
the hypothesis that low concentration olefinic hydrocarbons and nitric
oxide can greatly affect the rate of SCL photo-oxidation in air.
-------
-41-
Results of their experiments can be easily observed in Figure 2. The
rate constant obtained for the oxidation of SO in this experiment was
0.025 hr'1 with 40 ug/m3 (0.03 ppm) NO and .04 x 10"4 mol/m3 (.1 ppm)
cis-2-pentene present. They argue that due to reduced light intensity
in their chamber, their rate obtained would be analogous to a conversion
rate of about 10%/hour over London. Although this conversion rate seems
very high, they suggest it gives support to the oxidation rate of SO
of 0.65%/hour found in their earlier experiment as being more realistic
for existing atmospheric conditions than those found by Gerhard and
Johnstone,
Although the rate of S02 oxidation by ozone alone is quite slow,
Cox and Penkett (1971b) found that when SO was injected into a chamber
containing ozone and olefins the oxidation rate was greatly enhanced.
They also found that the oxidation rate, or aerosol formation, was
dependent upon the nature of the olefin; the rate of aerosol production
was much slower for terminally unsaturated propene and 4-methyl-1-pentem
than it was for the internally unsaturated olefins, cis-2-pentene and
3
2-methyl-2-butene. They calculated the oxidation rate of 270 ug/m
(.1 ppm) S0? in the presence of 2 x 10 mol/m (,05 ppm) ozone and
olefin to be approximately 3%/hr for cis-2-pentene and 0.4%/hr for proper-
As little as is known about the oxidation of SO in the troposphere,
still less is known about its oxidation in the stratosphere. Cadle and
Powers (1966) have suggested a possible 3-body reaction with atomic
oxygen
SO,, ••- 0 + M -* SO, -•• M
-------
-42-
1000.s
ro
E
x
a»
c
o
c
a>
o
c
o
o
CO
TJ
O
"5
«/}
O
i_
(D
O
<3-
O
CO
CVJ
T.
100.
10.
o S02 concentration
A Aerosol concentration
Inject
NO(.3xiO~4mol/m3)
2Pentene (.4l6X!0"4mol/m3)
I
100 200
Time(min.)
300
Figure 2. Aerosol formation and SO decay during the photooxidation of SO ,
-------
-43-
wherc M is a molecule of 0~ or N0, which acts to carry off excess energy,
thereby preventing prompt reversal of this reaction. Mulcahy, Steven,
and Ward (1967) found the rate constant for this reaction at room
temperature and from 0,7-3 torr (0.9-4x10 atra) to average (2.7±0,5)
56 -2 -1
x 10 m rnol sec „ Friend, Leifer and Trichon (1973) calculated
the rate of SO removal from the atmosphere by this reaction to proceed
quite slowly; 6 x 10" % hr~ at sea level and 2 5 x 10~4?6 hr"1 at 20 km.
Davis, Payne, and Stief (1972) suggested that the reaction of SO...
with the hydroperoxyl radical might also be of importance in the strato-
sphere
S00 -i- HO,, ->• S07 + OH
2. <- j
They found the rate constant for this reaction to equal 1,8 m mol
sec (t factor of 3). Based on estimated stratospheric H09 levels
made by Nicolet (1972), Friend et al. (1973) concluded that this
reaction would ultimately lead to the production of comparable if not
greater amounts of sulfate then would be produced by the reaction of
S09 with atomic oxygen
Friend et al„ (1973) also offer a reaction sequence showing the
3-bodied hydrolysis of S0_ to H SO and the subsequent hydrolysis of the
O £ *-T
sulfuric acid to H_SO. solutions, or acid embryos. These solutions
are neutralized by their reaction with ammonia to form "salt embryos".
The salt embryos act as nuclei for growing stratospheric particles on
which the oxidation of SO occurs as follows:
2SO + 2H 0 + 0 embry°> 2H SO (solution)
-------
-44-
Th s NH acts as a catalyst. When the ammonia is depleted from the
surrounding atmosphere, Friend et al. suggest that further absorption
and oxidation of SCL on the particle will be inhibited due to a drop
in pH as the buffer capacity of the NH. ion is exceeded. This reaction
sequence is thought to be responsible for the layer of HuSO. and sulfate
particles found in the lower stratosphere.
Perhaps the best way to summarize the possible reactions involving
SO in the atmosphere would be by repeating the summary made by Robinson
and Robbins (1968). "It seems that in the daytime and at low humidity,
photochemical reaction systems involving SO , NO and hydrocarbons are
of primary importance in t.ie transformation of SO into essentially an
H SO. aerosol. At night and under high humidity or fog conditions,
or during actual rain, it seems that a process involving the absorption
of SO by alkaline water droplets and a reaction to form S0~ within
the drop is a well-documented process and can occur at an appreciable
rate to remove SO from the atmosphere."
Environmental Sulfur Cycle
Table VII summarizes the estimates made by Eriksson (1960), Junge
(1963), Robinson and Robbins (1968), Kellogg et al. (1972) and Friend
(1973) in compiling their respective sulfur cycles. It should be kept
in mind that while all the values given are only estimates, the degree
of uncertainty in some is greater than that in others. For instance,
there are no measurements upon which estimates of the quantity of H?S
or S~ emitted by decaying land and sea biota can be based. Also,
there are no specifications of the form of sulfide emitted (i.e.,
H2S, HS~, or S~). These estimates have therefore been arrived at by the
-------
-45-
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-46-
authors' balancing of particular portions of their cycles; the difference
needed to balance each section has then been set equal to the flux of
il S or S from the land and sea to the atmosphere, respectively.
GLOBAL SULFUR CYCLE
A schematic global sulfur cycle is shown in Figure 3. This
diagram is a slightly modified version of that devised by Friend (1973).
The only modifications other than layout is the addition of atmospheric
oxidation schemes for H S and SCL, as well as a flux for the sorption of
SO by soils (Eriksson, 1963). The other fluxes shown in this diagram
are listed in column 5 of Table VII. All fluxes related to the
atmospheric sulfur cycle are discussed separately in the sections
concerning the sources and sinks of these compounds. The fluxes between
the pedosphere, hydrosphere,and lithosphere have been discussed in detail
by Friend (1973). All fluxes are given in Tg S/yr (10 g S/yr).
The sulfur inventory of the various reservoirs or spheres are
encircled and given in Tg S. All inventories are as shown by Friend (1973)
From cycles such as this the residence time of sulfur, or the time
taken for a complete turnover of the burden in a reservoir, can be
calculated for both the atmosphere and hydrosphere. The residence
time, T, of a substance is defined as T = M/R, where M is the burden or
inventory of the substance in a reservoir, and R is the total flux in-
to or out of that reservoir. The residence time of sulfur in the
atmosphere then is equal to
T = M/R =1.8 Tg/(217 Tg/yr) = 0.0083 yr = 2.7 days
-------
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-------
-48-
(Friend, 1973). The residence time of particular sulfur compounds would
be different. Hidy (1973) has calculated the atmospheric residence time
of SO to be 4 days and that of H S to be 2 days. Nordo (1973) has estimated
the residence time of SO to range between 20 minutes and 7 days. Individual
estimates of the residence time of sulfate in the atmosphere have not been
made, although it is probably comparable to that of sea salts, or about a
day or two (Eriksson, 1963).
Using Friend's data a calculation of the residence time of SO. in the
oceans can be made as follows:
T , M/R = .("±0.1) x 108Tg = (4_22±,324) X1Q6 years
3.08 x 10 Tg/yr
or about 4 million years. This estimate is a bit under half that of
Mackenzie and Garrels (1966). They calculated the residence time of S0~
in the oceans to be 10 million years based on river input data, Holser
and Kaplan (1966) have calculated the residence time at 21 million years
based on their own geochemical cycle. The values they use for the
sulfur inventory of the ocean and fluxes out of the ocean are essentially
the same as Friend's. However, they do not include an estimate of the
sulfur removed by marine plants. As a result, their fluxes are
underestimated by about 164 Tg/yr. Even without this flux though, the
residence time they calculated was not reproducible. The information given
in their cycle leads to a SO residence time of about 9 million years.
-------
-49-
CARBON CONTAINING GASES
CARBON MONOXIDE
Each year more carbon monoxide is released into the atmosphere than
any other pollutant (excluding carbon dioxide), and, each year the quantity
released increases. One would therefore expect a gradual increase in
ambient CO levels, yet one finds that the background concentration of
this gas in the atmosphere has not fluctuated the last few decades.
There must then be one or several major active sinks for CO within the
troposphere. Until just a few years ago though, investigations on
possible sinks had only turned up additional sources of CO, This dilemma,
as a result, came to be known as the "CO sink anomaly".
3
Background concentrations of CO range from 47-230 yg/m (0.04-0.20 ppra)
(Jaffe, 1973). Robinson and Robbins (1968) indicate that a mean concen-
3
tration of 100 yg/m (0.1 ppm) is found in the northern hemisphere,
while concentrations less than 58 yg/m (0.05 ppm) would probably be more
common in the southern hemisphere„
Sources
Carbon monoxide is the product, of incomplete combustion of fossil
fuels containing carbon. With the advent of large scale industrialization
and the tremendous increase in the use of the automobile, great quantities
of CO have been emitted into the atmosphere. Recent investigations for
CO sinks to explain why the ambient CO concentration has not been increas-
ing have resulted in the discovery of new natural sources of CO whose total
quantity far exceeds the total mass of CO produced as a result of man's
technology (Stevens et al., 1972).
-------
-50-
By far the greatest single anthropogenic source of CO is motor
vehicle exhaust, Jaffe (1973) estimated that of a total anthropogenic
q
CO emission source in the United States in 1970 of 132,6 x iO kg, 96.9
g
x 10 kg resulted from the burning of gasoline by motor /eludes alone.
Other significant contributions to this man-made CO burden are from
9
solid waste disposal (6.5 x 10 kg), industrial process loss (10.3 x
9 9
10 kg) and agricultural burning (12.5 x 10 kg).
Jaffe estimated the total CO emitted by combustion processes on a
9
global scale for 1970 to be approximately 360 x 10 kg, lie notes that
whereas the level of CO produced by man in the United Slates appears
to be leveling off, globally it is on an increase, Und;jr.ievoloped
nations, which are undergoing increased technological development are
not, as yet, concerned with the resulting pollution as muL.h as the
economic gains and so their emission levels are on the n ,e
The most widely recognized natural source of CO is forest fires
9
which have been estimated as releasing 11 x 10 kg CO into the atmosphere
each year (Robinson and Robbins, 1968),, This is, by no means, the only
major natural source of CO. Jaffe (1973) has written an excellent
review on all possible natural CO sources. The following is condensed
from that paper,
Minor amounts of CO have been found to be released from volcanoes
and marshes (Flury and Zernik, 1931). CO can also be formed during
electrical storms (White, 1932) and by the photodissociation of CO
in the upper atmosphere (Bates and Witherspoon, 1952); Calvert, Kerr,
Demerjian and McQuigg, (1972) have suggested the photodissociation of
formaldehyde as a possible source of CO and recently, Swinnerton,
Lamontagne and Linnenbom (1971) found CO to be present in rain water
-------
-5J-
in rather high concentrations. Galbally (1972) offered a hypothesis
wherein the CO in rain is a product of the photodecomposition of aldehydes
in the rain water by sunlight.
The ocean was first suggested as a major source of CO by Swinnerton,
Linnenbom and Lamontagne (1970), Linnenbom, Swinnerton and Lamontagne
q
(1973) have estimated the oceans can produce up to 220 x 10 kg each year,
q
whereas Liss and Slater (1974) have estimated this flux at 43 x 10 kg per
year. Robinson and Moser (1971) suggested that plants could indirectly
9
be the source of about 54 x 10 kg CO by the oxidation of released
terpenes. Finally, McConnell, McElroy and Wofsy (1971) suggested that
9
approximately 900 x 10 kg CO are produced each year by the oxidation
of methane.
In light of this new source information, Stevens et al. (1972) belici",
that natural sources of carbon monoxide could yield about 10 times more
CO than all anthropogenic sources in the northern hemisphere. Up to this
time it has been assumed that anthropogenic activity released far more
carbon monoxide than nature. It will be interesting to follow the out-
come of this contradiction over the next few years.
Removal Mechanisms
Carbon monoxide can be regarded, for all intents and purposes, as
being insoluble in water; its actual solubility being only 0,00234g/100g
HO at 20°C and 1 atm. Therefore, wet processes such as washout and
rainout can be regarded as playing an insignificant part in the removal
of CO from the atmosphere. Experiments on the absorption rate of CO by
an alfalfa canopy (Hill, 1971) showed that virtually no CO was absorbed
and therefore, vegetation can be disregarded as a sink. Absorption of
-------
-52-
CO by the oceans can now be disregarded as well because it has recently
been shown (Swinnerton et al., 1970) that the oceans actually constitute
a significant natural source of carbon monoxide. It seems then that
the major sinks for CO are gas-phase reactions in the troposphere and
stratosphere and absorption by soil fungi (Inman, Ingersoll and Levy,
1971, and Inman and Ingersoll, 1971).
There are two schools of thought concerning the sources and removal
mechanisms of CO in the atmosphere. The old school believed that anthro-
pogenic emissions far surpass natural emissions of CO into the environ-
ment, and, due to lack of information at that time, did not recognize
that chemical reactions in the stratosphere and troposphere play an
important part in destroying CO. Estimates of the residence time of
CO in the atmosphere in this school ranged from less than 4 years
(Bates and Witherspoon, 1952) to 2.7 years (Robinson and Robbins, 1968).
The new school recognizes the importance of natural sources of CO to the
tropospheric inventory as well as the importance of both stratospheric
and tropospheric reactions in removing it. Residence times estimated
by this school are on the order of 0.1-0,2 years (Weinstock, 1969;
Levy, 1971; Dimitriades and Whisman, 1971; and Levy, 1973b). In light
of the new information available on sources and sinks of CO, it may be
unfair to segregate residence time estimates into those arising from
old and new schools of thought. For clarification purposes the separa-
tion does seem necessary.
Soil: Experiments by Inman and Ingersoll (1971) showed that both
potting soils and natural spoils absorbed significant amounts of CO. They
observed that, in general, soils with the highest uptake activity were
-------
-53-
those with higher organic content and lower pH. Inman et al. (1971) also
noted that if a soil was autoclaved (sterilized) removal of CO by the
soil was inhibited. This suggested that the removal was due to biological
activity in the soil. Table VIII (Inman et al., 1971) summarizes the
results of tests on several natural soils performed in the laboratory.
Further tests isolated some aerobic micro-organisms which were present
in the soil and were then tested for their CO uptake abilities. The
following fungi were found to be active in removing CO from the soil:
4 strains of Penicillium digitatum, 1 strain of Penicillium restrictum,
4 strains of aspergillus, 1 of Mucor hiemalis, 2 of Haplosporangium
parium and 2 of Mortierella vesiculato.
In late 1972 Ingersoll published the results of a more extensive
study on the uptake of CO by soils. He measured the in situ uptake at
various locations throughout North America. Following is a short
summary of his findings.
1) Total amounts of CO destroyed by various soils ranged from
2
7.5 to 109.0 mg CO/m «hr,the spectrum ranging from tropical soils which
were the most active down to desert soils which were the least active.
Although there were many exceptions, more CO tended to be destroyed by
soils with low pH and moderate moisture content (^20%) than others.
2) The rate of CO uptake decreased as the concentration of CO
decreased in the air, with maximum removal with ambient concentrations of
100,000 yg/m3 (100 ppm).
3) Given the same soil, CO uptake was far greater when vegetation
was growing than when the soil was under cultivation. Ingersoll
suggests this is because the amount of organic matter present in soils
-------
-54-
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-------
-55-
being cultivated is substantially lower than that present in soils
actively growing crops.
4) Soils that were removed from their site of origin and tc^'
in the laboratory showed greatly reduced uptake ability. The magnu
of decrease was not uniform from one soil sample to the next^
Based upon data he had collected which he corrected for tempo* .
and uptake variations, Ingersoll (1972) estimated the CO uptake pu;
of the conterminous United States and the world. His estimates oi
9 12
505 x 10 kg and 14.3 x 10 kg CO per year for each, respective]},
3
based on ambient CO levels of 100,000 yg/m (100 ppm), three ordei•
magnitude greater than average ambient CO levels. Ingersoll no tec1
though that Seiler (197?) had found soil CO uptake rates to be one
of those Ingersoll had measured when these soils were exposed to atii:
CO levels (230-1200 yg/m ) rather than concentrations of 100,000 )ii;
On the basis of this information, Ingersoll reduced his estimates <•
total capacity of soils to consume CO in the United States and the.
9 12
to around 50 x 10 kg and 1.4 x 10 kg per year, respectively. ft.
Seller's article has not been published, it is impossible to state
whether his uptake measurements were taken in the field or the lab.
Ingersoll, as mentioned above, has found that uptake rates for thr
soil were greatly reduced when taken in the laboratory as compared
field measurements. Unless measurements were taken by both Ingerso•
and Seiler under the same conditions, one must question the valid!i
in using Seiler's data to reevaluate estimates made using Ingersol'
data. Since the estimates made at this stage are only a first
approximation, perhaps such concerns are unimportant. Earlier, In,
-------
-56-
Ingersoll (1971) also estimated that soils of the continental United
g
States were capable of removing over 500 x 10 kg CO per year.
Based on measurements of the CO uptake rate of two soils, Heichel
(1973) concluded that the resistance of the soil to CO removal increases
as the moisture content of the soil decreases, and also increases as the
degree of air circulation above the soil decreases. Comparing his
uptake and resistance measurements with those of Turner, Rich and
Waggoner (1973) who measured the uptake rate of ozone on one of the same
soils, Heichel found the rate of CO removal by the soil to be 2-3 orders
of magnitude less than the removal of 0 . There was also a correspond-
ingly greater resistance to removal of CO than 0 . Heichel determined
on the basis of his mt-asurements, and assuming an average ambient CO
level of 350 yg/m (0.3 ppm), that the land can remove the equivalent
of 14% of the tropospher:.c CO inventory of 4.75 x 10 kg CO (Weinstock
g
and Niki, 1972) each year, or about 67 x 10 kg/yr.
Recent experiments by Smith et al. (1973) also showed that soils
are capable of effectively removing CO from the atmosphere. They also
found though that when mcist soils were placed in the chamber with air
containing 100,000 yg CO/m (100 ppm), the concentration of CO in the
chamber rose, in one case tripled, before it was reduced to zero. This
effect was found to be more pronounced in sterilized soils than
unsterilized soils. They surmised that the evolution of CO from the
soils is a nonbiological process, but made no attempts to identify the
processes responsible. Seiler and Junge (1970) also noticed this effect.
Their test results showed that at ambient CO levels (200 ug/m } or 0.2 ppm)
the rate of CO production by the soil is balanced by the rate of CO
-------
-57-
absorption by the soil. This finding lends another complication to the
identification of sources and sinks for CO. Perhaps ultimately it will
be shown that soils do constitute a major sink for carbon monoxide; first
we need to know if there is evidence of CO evolution from in situ soils
or whether this is a phenomenon only found in laboratory experiments,
and we need quantitative data to see what the ratio of CO evolved to CO
removed from the atmosphere is at different ambient CO levels.
Atmospheric Reactions Involving CO: Recently it has become
apparent that the stratosphere constitutes a sink for carbon monoxide.
The significant reaction seems to be the reaction of CO with the hydroxyl
radical (Weinstock, 1959; Pressman and Warneck, 1970; and Pressman et al.,
1970) as follows:
OH + CO -> CO + H (1)
The rate constant for this reaction at -56°C, the temperature of the
4 3
lower stratosphere is k = 6.02 x 10 m /mol*sec (Schofield, 1967).
Westenberg (1972) found that the rate constant increases to 10.2 x 10
3
m /mol-sec at room temperature (27°C) . Reaction 1 is followed by the
sequence:
H + 02 + M -> H02 + M (2)
CO + HO -»- CO + OH (3)
L* £.
The rate constant for reaction 2 has been found to be equal to 4.35 x
O 7 A 9
10" m /mol -sec (Westenberg and de Haas, 1972).
-------
-58-
The significance of reaction 3 has recently been the subject of much
debate and experimental work. It was previously believed that the regen-
eration of the OH radical proceeded via the reaction
H + H02 •* 2 OH (4)
Davis et al. (1973) have concluded that reaction 3 proceeds quite slowly
3 3
(k probably being less taan 6 x 10 m /mol'sec) and therefore, should not
play a significant role in atmospheric chemistry. Westenberg though
believes that reaction 3 probably proceeds even faster than reaction 1.
Westenberg and de Haas (1972) derived a rate constant for reaction 3 at
room temperature of about 6.02 x 10 m /mol*sec which is indeed faster
than that of reaction 1. Explanations offered to explain the difference
in these k values (Davis, Payne and Stief, 1973) are not conclusive.
O
It has been pointed out (K. L. Demerjian, 1974) that the rate constants
for reaction 3 are erroneous. For further discussion of these reactions
and many more of interest the reader is referred to a recent work by
Demerjian (1973).
Naturally, if this mechanism is correct it would require that a
significant amount of hydroxyl radicals be present. Assuming that CO
is in a steady-state in the atmosphere with a concentration of 100 yg/m
(.1 ppm) Weinstock (1969) calculated the OH required to maintain this
steady state condition according to the equation:
[OH] = P/k[CO]
where P = production rate of CO and k is the second order rate constant.
-3 3
He calculated that 1.16 x 10 mol OH/m were required and that requirement
could easily be satisfied by an OH regenerative process as suggested above.
-------
-59-
Vertical CO profiles of the atmosphere have recently been measured
by Seller and Junge (1969), Seller and Warneck (1972) and Goldman, Murcray,
Murcray, Williams, Brooks and Bradford (1973) . These first two reports
both found a sharp decrease in the CO mixing ratio when crossing the
tropopause into the stratosphere. The mixing ratio is defined as the
mass of gas per unit mass of dry air. Goldman et al's. profile did not
show a sharp decrease in the mixing ratio at the tropopause but they
believe this might be due to local atmospheric conditions just around
Holloman, New Mexico, where their measurements were taken. It seems
obvious from vertical profiles such as these that there is a significant
decrease in CO levels above the tropopause, thus supporting the conclusion
that the stratosphere is a sink for carbon monoxide.
Quantitative estimates of the CO destroyed in the stratosphere by
this reaction mechanism have not been made, although Pressman and
Warneck (1970) believe that virtually all CO entering the stratosphere
is so destroyed. Therefore, the size of the stratospheric sink is
dependent upon the transport rate of CO rich air through the tropopause
into the stratosphere. They have estimated this flux at 1.3 x 10
2
mol/m *sec but admit the degree of uncertainty in this estimate is very
high due to insufficient data. Based on this flux and recent estimates
of the total CO reservoir in the troposphere, they estimated that about
11-15% of the total CO inventory in the troposphere is destroyed in the
9
stratosphere by the reaction sequence mentioned above, or about 52-71 x 10
kg/yr.
Although the above reaction sequence has been recognized as a major
sink in the stratosphere, only recently has it been suggested to be a
-------
-60-
sigrtif leant destructive reaction in the troposphere (Weinstock, 1969) .
Recent investigations (Levy, 1971; McConnell, McElroy and Wofsy, 1971)
have shown that sufficient amounts of hydroxyl and hydroperoxyl radicals
exist in the troposphere to provide a mechanism for CO oxidation there
as well .
Levy (1971) suggested the following reaction sequence which provides
for both the production of hydroxyl and hydroperoxyl radicals and the
oxidation of CO and NO in the troposphere.
(1) 0 + hv (2900A • 0('5P) + M
73 -1 -1
k = 3.01 x 10 m mol sec
(3) OD) + HO -> :>OH
k = 1.8 x 10 m mol sec
(4) OH + 03 -> H02
53 -1 -1
k £ 3.01 x 10 m mol sec
(5) OH + CO •+ C02 + H
43 -1 -1
p = 9.0 x 10 m mol sec
(6) H + 0 + M ->- HO + M
k, = 1.08 x 104 m6 mol"2 sec 1
D
(7) HO + NO -*• NO f OH
^ z
k = 3.01 x 10 m mol sec
More recent estimates of these rate constants by Kummler and Bauer (1973)
are in excellent agreement with those of Levy listed above.
-------
-61-
Although quantitative estimates of the total CO possibly consumed by
this reaction series have not been made, Levy did calculate a carbon
monoxide lifetime in the surface atmosphere of 0.2 years using hydroxyl
radical concentrations which he had determined. This estimate is in good
agreement with that made by Dimitriades and Whisman (1971), but is still
the subject of considerable dispute among other scientists (Kummler,
Bortner and Jaffe, 1971) who produce their own evidence to show that gas
phase kinetics at ground level cannot alone explain the sink anomaly.
Carbon Monoxide Cycle
The last several years have seen an explosion of interest in
defining natural CO sources and sinks. While several researchers have
confirmed that soil and the reaction of CO with hydroxyl and hydroperoxyl
radicals in the stratosphere and possibly the troposphere constitute
a major sink for CO, other researchers have found additional natural
sources of this gas. These sources, which include oxidation of methane,
oxidation of terpenes and the oceans to name just a few, are now thought
to contribute more CO to the atmosphere than that emitted from anthropo-
genic activity. This is an outright contradiction of what only two years
ago was thought to be the final word; i.e., anthropogenic CO emissions
are many times greater than natural CO emissions.
Table IX compares recent estimates of the importance of CO sources
and sinks. It is obvious from this table, if these estimates are at all
reasonable, that soil and gas-phase oxidation in the stratosphere and
troposphere might be capable of consuming all the CO that nature and man
produce. If so, this explains why the background concentration of CO
has not increased over the last several decades, and we might suggest
that there no longer exists a "CO sink anomaly".
-------
-62-
Table IX. ATMOSPHERIC FLUXES OF CARBON MONOXIDE (in 10 kg/yr)
I. Sources
a. anthropogenic
b. natural
1. oceans
359
43-220
2. oxidation of terpenes 54
3. oxidation of methane 900
4. other 1_
TOTAL 1356-1533
Jaffe (1973)
Liss and Slater (1974)-
Linnenbom et al.(1973)
Robinson and Moser (1971)
McConnell et al. (1971)
II. Sinks
a. soil 67-1400
b. gas-phase oxidation
1. stratosphere 52-71
2. troposphere ?_
TOTAL 119-1471
Heichel (1973)-
Ingersoll (1972)
Pressman and Warneck (1970)
-------
-63-
NITROGEN CONTAINING GASES
Although the most abundant oxide of nitrogen in the lower atmosphen
is nitrous oxide (N«0), it does not play an important part in air pollu
tion chemistry. The two nitrogen oxides which are important are nitric
oxide (NO) and nitrogen dioxide (NO ) for they play a prominent role in
the generation of photochemical smog. In fact, the presence of nitrogen
dioxide, as a light energy acceptor, is a necessary condition for
photochemical reactions to proceed.
Measurements of the background levels of nitrous oxide consistently
average to about 460-490 yg/m (0.25-0.27 ppm) (Schutz, Junge, Breck
and Albrecht, 1970). Ambient N09 levels are much lower, probably around
3 3
2 yg/m (.001 ppm) (Robinson and Robbins, 1968) to 2.6 yg/m (.014 ppm)
(Junge, 1963). Too few accurate measurements have been made of ambient
levels to obtain a meaningful average, but Lodge and Pate (1966) found
concentrations up to 11 yg/m (.006 ppm) in Panama. Not many measure-
ments of NO levels have been made but they are approximately the same
as those of NO .
The other nitrogen compound considered is ammonia. Although ammon ;.<"
per se, is a relatively unimportant air pollutant, it does play an
important role in atmospheric chemistry through its part in the formation
of aerosols. Ambient concentrations of ammonia seem to average about
4 yg/m (.006 ppm) (Robinson and Robbins, 1968).
-------
-64-
NTTROGEN OXIDES
Sources
Nitrous oxide (N-0.), the most abundant nitrogen compound in the
atmosphere, is produced from bacterial decomposition of other nitrogen
compounds within the soil. Arnold (1954) was the first to study the
production of N00 by soils. He found that the factor most conducive
t~
to increased NO evolution is a high soil moisture content, especially
if a nitrogen source, stch as nitrate or ammonia, is present. Bacterial
decay is the primary source of nitrous oxide and so is probably responsible
for ambient N?0 concentrations. An estimate of the flux of N90 from
<-> ^
soils into the atmosphere has recently been made by McConnell (1973).
He estimates that although an unknown amount is released and reabsorbed,
about 1.1 x 1010 kg N 0 - N (1.73 x 1010 kg N20) are available for oxida-
tion or photolysis to NO and N in the stratosphere each year. Schutz
.\ L»
et al. (1970) declined to extrapolate the data they collected on three
soil samples to attain a global flux. Their measurements showed a flux
Q 7
on the order of 10 g N/;)0/m 'sec, an order of magnitude which, if
maintained globally, would necessitate an NO residence time of about
70 years. This is the same residence time they estimated based on the
photodissociation rates of N~0 in the troposphere and stratosphere.
Goody and Walshaw (1953) estimated a globalN?0 production rate of about
12
100 x 10 kg/year and Robinson and Robbins (1968) suggested that soils
10
produce about 59.2 x 10 kg NO each year by biological action; and
of this about 55.4 x 10 kg (35.3 x 10 kg NO - N) are reabsorbed
by the soil and about 3.8 x 1010 kg N20 (2.4 x 1010 kg N20 - N) travels
up to the stratosphere where it is destroyed. It is this latter rate
that is shown in Table XI.
-------
-65-
Craig and Gordon (1963) raised the possibility that the ocean
might be a source or a sink of NLO. Bates and Hays (1967) concluded
though that the uptake of N^O by the oceans in areas where upwelling
waters are deficient; in NO is of negligible importance. Based upon the
difference between the mean N_0 concentrations of oceanic surface waters
and adjacent air, and an estimate of the total liquid phase resistance,
Liss and Slater (1974) estimated the flux of N?0 from the ocean to the
atmosphere at 1.2 x 10 kg/year. This flux is based entirely on theory,
but indicates that the oceans could conceivably release significant
quantities of N?0 to the atmosphere. Laboratory and field experiments
are needed to decide whether or not the ocean is a source of N?0.
Production rates of NO and NO by soils are much more difficult to
measure or estimate, and good data are lacking. McConnell (1973)
recently summarized a few of the problems involved in appraising the
amount of nitrogen oxides produced by soil. He contends that this soil
source is small compared to that produced as a result of the gas phase
oxidation of atmospheric ammonia by OH which he thinks produces 7 x 10
kg NOY - N per year. McConnell offers alternative reaction sequences
A
for NH in the atmosphere; one reaction sequence provides a constant
source of NO, the other a sink. If the latter is shown to occur in
the atmosphere, in order to account for the amount of NO known to be in
the atmosphere, then an additional source of NO must be found. In this
case, McConnell concedes that the soil might actually constitute a
significant source of NO , on the order of 10 kg/yr.
A.
Nitric oxide is also produced as a result of the photolysis of
N_0 in the stratosphere as follows:
-------
-66-
N'20 + hv -»• NO + N
-9
The photolysis rate for this reaction is less than or equal to 7.4 x 10
sec (McElroy and McConaell, 1971). Whereas these authors estimate the
o
production of about 3 x 10 kg NOY - N per year in this manner, Bates
A
and Hays (1967) estimated that 3.5 x 10 kg NO are produced annually
by this photolysis reaction.
The other primary source of nitrogen oxides is anthropogenic,
primarily combustion processes. Estimates of production rates for NO
and N0? are included together because available emission data rarely
distinguish between these two forms. Robinson and Robbins (1970)
" 9
estimated that 53 x 10 kg of NO , or 16 x 10 kg NO- - N were emitted
annually (here again NO includes both NO and N0? production).
I_n toto, natural emissions of NO (including NO) are approximately
9 9
15 times greater than anthropogenic emissions (768 x 10 kg vs. 53 x 10
kg N00) (Robinson and Robbins,, 1970). Therefore, anthropogenic emissions
play only a minor part in the total circulation of nitrogen compounds
in the atmosphere.
Removal Mechanisms
At 20°C and 1 atm total pressure, the solubility of NO in water
is 0.121 g/lOOg HO. Upon release from soils, an unknown portion is
believed to be removed by vegetation, soil and water. No information is
available on these reactions. The major portion of the released N?0
is destroyed by photodissociation in the stratosphere and upper tropo-
sphere. Under normal tropospheric conditions, N-0 is chemically inert
and it partakes in no other chemical reactions. The residence time of
-------
-67-
nitrous oxide is probably around 70 years if the,r-
biosphere (Robinson and Robbins, 1968), but coul'i
years if there is a biologic loss mechanism. Hid, i i'i
residence time as 4 years.
Nitric oxide is rather insoluble in water. >!:
pressure, its solubility is 0.00618 g/lOOg HO •'
the other hand, immediately dissociates in walti •
For this reason the solubility of N0_ in watei i - • :•
Nitric oxide is either oxidized to NO or [/
is then removed primarily be precipitation, more ••' • ••
form of nitric acid (HNO ). It can also be abbu^ ;
soils or participate in photochemical reactions : »•
their reactivity, the residence timesof NO and '••' .••••
probably around 5 days (Hidy, 1973) . Nitrogen *' . >•• "
the atmosphere by the following mechanisms.
Vegetation: Vegetation has been shown cap.;!"'' >>!
cant amounts of NO and NO from the atmosphere, i; •.<
that alfalfa and oats absorbed NO from the aij
2
mol/m -sec when exposed to an atmosphere contain •>••• !;-
-5 3
M x 10 mol/m ). With time, the uptake rate u- <*•
to air containing between 300 and 460 yg NO,,/in '•• " •
uptake rate remained constant when the plants v/u-' • •
150 ug NCL/m . He also found that there was nn >'( < • • •
rates when the intensity of the sunlight was ivi • "';.
Tingey extrapolated his data to estimate !L- ;-.•""•
the Salt Lake Valley in Utah. His estimates, iv-.-; -«
-------
-68-
2
of about 2000 kin in the valley with an 85% vegetal covering in the summer,
are tabulated below (Table X).
Table X. WTES OF NO UPTAKE IN SALT LAKE VALLEY, UTAH
Removal Rate „ , _.
,, „ Removal Rate
NO^ Cone. ,2 ^ ,, 11
2 per km from Valley
(24-hr average) (kg/day) (kg/day)
40 yg/m3 18 36 x 103
20 yg/m3 9 18 x 10
10 yg/m3 4.5 9 x 103
On an annual basis, removal of NO from the valley with an ambient
NO concentration of 9.6 yg/m (about 3f-4 times greater than background
NO levels) would total about 3 x 10 kg N0?. This may be compared with
the estimated total global anthropogenic annual N0? emissions of
53 x 109 kg (Robinson and Robbins, 1970).
Hill (1971) found in his experiments on the uptake rate of gases by
an alfalfa canopy that NO was absorbed with a depositional velocity of
0.1 cm/sec. NO was absorbed at a velocity of 2 cm/sec when present
in the air of the chamber at a concentration of 2 x 10 mol/m (or
96 yg/m ). Using ambient N02 concentrations found in thbse areas of
Southern California from August to October of 1968, and assuming a
continuous alfalfa cover, Hill estimated that N02 could be removed at
T
a rate of 0.1 g/m^'day. Because nitrogen is often a limiting factor in
-------
-69-
plant growth, Hill believes his study shows that NO, and ;j'i rc
as much as .03 g N/m -day. Since NO dissociates in walei .i
that the N absorbed from the atmosphere as NO would be me I •''••
so the amount absorbed over one year could be significant
Soil: Nitrogen oxides (especially N_0) have long be- ; I
"'"" Z*
produced by biological action in soils. Recently though,
(1971) found that soils could absorb nitrogen dioxide fivn
as well. They found that when they passed air containing
in a test chamber, the concentration of NO in the air wa., .
3 3
an initial value of 190 x 10 yg/m (100 ppm) down to S, ' !
(3 ppm) over a 24 hour period. When the soil was autoci;
NO-present over the same time period was reduced from J8( '
(97 ppm) to only 25 x 10 yg/m (13 ppm) . The reaction l.\ '• <
be the cause of the NO- uptake in the soil was not discovt
Extrapolating the results of this experiment, the author.:, •-•,
the soils of the United States might be capable of remcn i; -
of NO- per year from the atmosphere, an amount they point ••>•''-
bit under 20 times the total annual production of NO, 311 i i-: •
(3.3 x 10 kg). Of course, such a removal rate is larg'-i
for there are too many variables that need to be taken ui: • .
such as soil moisture content, soil composition, both iiuix,",.:1
organic, and vegetal covering to name just a few. AithouHi T
such an estimate is warranted, the limitations placed upon > 'i
must be kept in mind until more information becomes avajJ, :M
Nelson and Bremner (1970) point out that the NO ahst . -
will ultimately be oxidized to nitrate. These nitrates c < <
-------
-70-
decompose and result in the production of nitrogen dioxide again. The
-4
rate of N0? production by nitrate breakdown in soils is 2 x 10 g NO /
9
m"*hr (Marchesani, Towers, and Wohlers, 1970, and Makarov, 1970, as cited
in Bohn, 1972). This NO production rate is dependent upon the nitrate
content of the soil and does not proceed during darkness.
Nitric oxide may also be absorbed by soils, but is then oxidized
almost immediately to NO (Mortland, 1965; and Bremner and Nelson,
1968). Mortland has also discovered that transition metal ions in
the soil promote NO absorption. If the soil is saturated with alkaline
earth cations though, absorption of NO is halted. Sundareson, Harding,
May, and Henrickson (1967) found that alkaline-earth zeolites readily
absorb NO and release it as NOY and HNO_ when heated. To date, the
A o
role organic matter plays in the absorption of nitrogen oxides by soil
remains a mystery. Ganz, Kuznetsov, Shlifer, and Leiken (1968, as cited
in Bohn) found that upon passing NOY-contaminated air through 1 meter of
A
peat,all nitrogen components were removed. Organic matter is such an
important component of soil that to not be fully aware of its effects
on a gas could only hinder understanding of the mechanism of absorption
by the soil. Obviously, nore research is needed in this area.
Water Bodies: There are very little data available on the amount of
nitrogen oxides absorbed by the oceans. Craig and Gordon (1963) first
suggested that the oceans might constitute a sink for N_0 when they found
that the sea water at depth was depleted in N?0 compared to the surface
waters which they found tc be in equilibrium with atmospheric concentra-
tions. This depletion could not be explained by temperature variations
with depth. Based on the mean upwelling speed of the ocean's waters,
-------
-71-
-4 -5
thought by Bowden (1965) to range from 10 to 10 cm/sec, Bates and
Hayes (1967) estimated the potential sink strength of the oceans for N?0
-7 6 ?
as ranging between 6 x 10 and 6 x 10~ kg/m -yr- • Ljss and Slater
(1974), on the other hand, concluded that the flux of N90 uas from the sea
to the air. They calculated a flux rate of 3.2 x 10 kg/m""yr based upon
an N«0 concentration gradient across the air-sea interface measured by
Junge and Hahn (1971) . The total flux of NO from the sea to the
atmosphere, if assumed constant over the entire oceanic surface, is
1.2 x 10 kg/yr. With such an obvious contradiction as to the direction
of the NO flux, there is a real need for additional research in this area.
Washout and Rainout: The major sink identified thus far for nitrogen
oxides (NOY) is the solution of soluble species in cloud anil rain droplets
A
with subsequent removal by precipitation. According to Haagen-Smit and
Wayne (1968) the reaction is:
4NO
Georgii (1963) believes the reaction proceeds as follows:
2NO + HO -v HNO + HNO
£• £. O .£.
3
Georgii found when a total volume of 20 m of atmospheric air containing
average concentrations of NO , but cleansed of all natural aerosol
particles, was passed over both distilled and rain water at a velocity
of 5-10 cm/sec, that the nitrate (NO ) concentration in the rain water
O
increased from 0.18 to 0.45 mg/1 while the nitrite (N0~) concentration
increased from 0.04 to 0.08 mg/1. Georgii further showed that the N0~
3
concentration in rain is related to the concentration of NU in
-------
-72-
the atmosphere. In Frankfurt am Main, Germany, where ambient NO levels
are higher in winter than summer, Georgii found that the nitrate content
in rainwater was also greater in winter than summer.
Robinson and Robbins (1968) suggested that the dissolution of NO
proceeds as follows:
3NO,, + HO + 2HNO + NO
£* £. J
However the hydrolysis reaction proceeds, the outcome is the same in
all cases; the nitric acid formed is absorbed onto hygroscopic particles
or reacts with atmospheric ammonia to form nitrate salt aerosols (NH NO
for instance). It is thea either removed by precipitation, or if
vaporization of the drool3t occurs, by dry deposition.
McConnell (1973) estimates that 2 x 10 kg NO" - N is removed from
the atmosphere each year by precipitation, and an additional 7 x 10 kg
NOY - N removed by dry deposition, the major portion of this probably
A
being HNO .
Atmospheric Reactions Involving Nitrogen Oxides: Although more
nitrous oxide is released to the atmosphere than any other nitrogen
oxide it does not play a major or very complex role in atmospheric
reactions. Because it is chemically inert in the troposphere, its sink
lies in the stratosphere where it is transported by vertical mixing and
destroyed by photolyzing reactions. Bates and Hays (1967) indicate
that the most significant reactions are:
N20 -- hv ->- N2 + 0 (1D) ; A < 3370 A
N20 i- hv •> NO + N (4S) ; A < 2500 A
-------
-73-
The latter reaction they consider responsible for about 20% of the total
dissociation in the stratosphere. They calculate an average latitudinal
-9 -1
dissociation rate of about 3 x 10 sec
Nitric oxide can be removed from the atmosphere by several reactions.
The primary reaction is its oxidation by ozone to form NO
NO + 03 + N02 + 02
18 ^ 1
which has a rate constant of 9.0 x 10" exp (-1200/T) m sec" (Schofield,
1967). In the upper atmosphere, NO can be photolyzed to atomic N which
may react with other NO molecules to form N as follows:
NO + hv -> N ( S) + 0 ( P) ; photolysis rate = 4.0 x 10 sec (Callear
and Pilling, 1970)
N + NO + N. + 0;k = 2.2x 10~U cm3 sec"1 (Schofield, 1967)
Nitrogen dioxide also can engage in a number of reactions. It may
be oxidized to gaseous NO, by the reaction
«5
N02 + 03 + N03 + 02
•JO -7
The rate constant equals 9.0 x 10" exp (-3500/T) m /sec (Schofield, 1967).
Or, it may form nitric acid by its reaction with hydroxyl radicals as
follows:
OH + N02 j HON02* -> HN03
The nitric acid formed is removed by precipitation. For further elaboration,
the reader is referred to McConnell and McElroy's (1973) article.
The importance of NO and NO as pollutants reflect their participation
in photochemical reactions. In polluted atmospheres they react with SO-
-------
-74-
and hydrocarbons to forn aerosols. Probably the most important photo-
chemical reaction involving NCL is its photodissociation as follows:
N02 + hv (2900 A < A < 3800 A) t NO + 0
This atomic oxygen then is free to react with molecular oxygen and form
ozone.
Peroxyl radicals (ROO)> formed between reactive free radicals (R)
and CL, can react with NO and NO to form alkyl nitrates or peroxyacyl
nitrates as follows:
ROO* + NO ->• ROONO- h^ R0« + NO
ROO' + NO -»• ROONO
^- L*
These secondary reaction products are targets of further photochemical
attack. For example, peroxyacyl nitrate may be photodecomposed into NO
and acylate radical as follows:
0 0
11 hv "
RC-OONO + RCO- t N02
This rather brief overview of possible photochemical reactions involving
nitrogen oxides is abridged from Haagen-Smit and Wayne (1968) . It is
inappropriate here to elaborate on photochemical reactions involving
nitrogen oxides. The reader is referred to Altshuller and Bufalini (1971),
Cadle and Allen (1970), and Leighton (1961) for more detail.
Environmental NO Cycle
A
The circulation of nitrogen oxides in the atmosphere is complex
and not well understood. The global NOY cycles formulated by Robinson
-------
-75-
and Robbins (1968, 1970) and McConnell (1973) make clear how very little
good quantitative data exists. A more complete global nitrogen cycle
(Figure 4) for both nitrogen oxides and ammonia is presented at the end
of the following section on ammonia.
The holes in these cycles are obvious and major ones (see Table XI).
We need, for instance, a better idea of how much NO and NO,, is released
from the soil, how much NH is oxidized to NOY, and how much NO and N00
J A /
is destroyed by photochemical reactions. We also need to know why large
discrepancies exist between estimates, and the bases for these estimates.
For example, McConnell bases his estimate of NO removed by dry deposi-
A
tion on soil data, the nitrogen content of precipitation and a deposition
velocity factor. Robinson and Robbins, on the other hand, use the
deposition velocity function to estimate gaseous deposition, a removal
mechanism McConnell makes no mention of. These and other inconsistencies
indicate the need for further study.
-------
-76-
10
Table XI, ATMOSPHERIC FLUXES INVOLVED IN VARIOUS NOY CYCLES (10 kg N/yr)
A
I, Sources
a. anthropogenic; NO, NO
b. biological; N?0
c. biological; NO, N0?
d. oxidation of NH
e. stratospheric transport;
NO, N02
TOTAL
11. Sinks
a. rainout
b. dry deposition
c. oxidation of N20-»NO (strat.)
d. photolysis of N20->N2 (strat .)
e. gaseous deposition
TOTAL
Robinson and
Robbins
(1968)
1.5
1.2
30.4
N.E.
33.1
112.9
27.1
0.2
10.7
150.9
Robinson and
Robbins
(1970)
1.6
1.2
23.4
N.E.
26.2
7.5
1.9
0.2
4.5
14.1
McConnell
(1973)
1.8
1.1
?
7
0.07
9.97
2
7
0.03
(1.07)
(10.1)
N.E. = Mechanism recognized but no estimate made.
*NOTE: Friend (1973) has pointed out that due to an error in converting units
of kilograms per hectare to tons per square meter, much of Robinson and
Robbing nitrogen compound cycle is invalid. Friend though, gives no
indication as to which values are wrong. Due to lack of information,
their cycle is nonetheless shown for comparison.
-------
-77-
AMMONIA
Sources
Most ammonia in the atmosphere is the result of the bacterial
decomposition of organic material on the earth's surface. The factors
which affect the emission of this NH from the soil are its nitrogen
O
content, pH, and moisture content (McConnell, 1973; and Georgii, 1963).
NH is more readily released from dry soils than moist ones, and is more
readily released when the pH of the soil is greater than or equal to 6.
It has been suggested that in acid soils NH. is probably biologically
oxidized (autotrophic nitrification) to NO and therefore not available
ij
for release. The net oxidation reaction would be
NH* + 202+ NQ~ + H20 + 2H+
(Keeney, 1973). DuPlessis and Kroontje (1964), on the other hand,
suggested that the low hydroxyl ion activity in acid soils and adsorption
of NH on soil colloids prevents or retards the release of ML, according
to the reaction
NH* -H OH~^r NH3 -*- HO
In alkaline soils the activity of OH is greater than that of H and so
NH is readily released.
Junge (1963) suggested that the oceans may contribute some NH to
the atmosphere, but to date there has been no accurate measurement made
upon which to base an estimate of the source strength. McConnell has
estimated that these biological sources release about 17 x 10 kg NH -
N each year, whereas Robinson and Robbins ' (1968) nitrogen cycle calls
-------
-78-
12
for the release of NH, on the order of 10 kg NH_ each year (see
Table XV),
Anthropogenic NH emissions result primarily from the combustion of
O
coal. Robinson and Robbins (1970) estimated that the atmospheric ammonia
burden due to man's activities is 0.4 x 10 kg NH» - N per year, less
«D
than 2 1/2% of the estimated NH burden due to biological emissions
(17 x 1010 kg).
Removal Mechanisms
Ammonia is extremely soluble in water. At 20°C and 1 atm its
solubility in water is 62.9 g/lOOg H?0. Ammonia therefore can be
readily absorbed by water bodies and vegetation. It is also removed
by its reaction with hydroxyl radicals to form nitric oxide. Probably
the most important removal mechanism though is its solution in rain
water along with S0~ or other gases to form aerosols such as (NH,)7SO..
The quantity of NH removed by precipitation is ordinarily determined by
*J
measuring the quantity of NH. present in rain. From the data available,
the residence time of ammonia in the atmosphere is probably around 7 days.
The following mechanisms are important in removing ammonia from the
atmosphere.
Vegetation: Because it is normally present in the atmosphere in
such small concentrations, ammonia's interaction with soil, water and
plants has long been ignored. Hutchinson, Millington and Peters (1972)
and Porter, Viets and Hutchinson (1972) have found that even at the low
naturally occurring atmospheric concentrations, plant leaves absorb
significant quantities of ammonia. Porter et al . exposed air containing
-------
-79-
N labeled NH to corn seedlings in growth chambers. '!'!•<' -<.<\i root
O
system was isolated from the atmosphere by polyethylene bup i?'i>ch were
tested and found to be impermeable to NH . At the condition of" the testing
O
period, the plants were analyzed and found to contain "N ,v-i "."od nitrogen
compounds. They therefore deduced that direct foliar absorp1 i--n was
responsible for the presence of most of the N found in f.h-- :>' nt.
They have indicated that simple isotopic exchange might a<:,c»;iut tor a
small part of the apparent [ N] NH uptake. They did not att'.-mpt to
establish an uptake rate for NH by these plants.
Hutchinson et al. (1972) based their estimates of Id1,' u|.t >K; rates
of various plants on the difference in the mass flows oF iv|i ,u-'. ,i-..js-ed
absorption rates of NH showed large diurnal fluctuation- . liii/- « '••[< --uecies
and for soybeans at three different nitrogen fertility lev."
Soil: The absorption of NH by soils has received rvi.'i rv.'y little
attention until recent years. Malo and Purvis (1964) ijiV".i ii;i'fd the
seasonal variations in atmospheric NH concentrations over K'.1^ Jersey in
O
relation to the rate of NH absorption by six New Jersey '.^"is. The
•J
atmospheric NH content averaged 57 yg/m , considerably -i>(.•••[<>-< i ban the
normal background concentration of about 4.2 yg/m . The fM. v..-. dtmospherJc
-------
-80-
Table XII. AMMONIA ABSORPTION RATES FOR FOUR CROP SPECIES AND FOR SOYBEAN
AT THREE NITROGEN FERTILITY LEVELS
Added
„ . soil
Species
^ nitrogen
Soybean (Glycine max)
Soybean (Glycine max)
Soybean (Glycine max)
Sunflower (Helianthus annuus.)
Corn (Zea mays)
Cotton (Gossypium hir^utum)
(mg)
0
5
20
l 5
5
5
Leaf
area
(cm )
65
80
84
96
58
55
NH3 in
chamber air
(yg/m3)
29
24
24
31
24
44
NH3
uptake
rate,
(yg/m -hr)
410.
420.
400.
490.
560.
350.
NH_ in New Jersey is probably due to the combustion of coal, oil and
J
gasoline. The rates of absorption by six soils under field conditions
in New Brunswick are given in Table XIII, along with some soil character-
istics.
The factors governing the amount of NH absorbed by soils were
O
investigated by Hanawalt (1969a,b). Hanawalt (1969a) tested the
influence of several atmospheric factors on the sorption rate which
included the length of exposure, the ambient NH, concentration, the
ambient air temperature and the velocity of the air passing over the
soil. The most important factor causing a change in sorption rates was
the ambient NH concentration. As the concentration increased from
O
3 3
about 37 yg NH - N/m to about 67 yg NH - N/m , the sorption rates
J J
-------
Table XIII,
-81-
SOIL CHARACTER AND ABSORPTION RAThS
FIELD CONDITIONS
Date and Soil
November 21, 1962
1. Collington sandy loam
2, Dutchess loam
3. Lakewood sand
February 18, 1963
4. Matapeake silt Ica/i
5. Dutchess shale loam*
6. Nixon loam*
PH
4.6
5.4
4.0
5.6
7.0
5.7
Organic
matter
2.32
7.12
2.51
2.92
--
q
0
li
26
0
Ifa
-
_
I
I
O.t. :
*Soils 5 and 6 both from cultivated fields
**Cation Exchange Capacity (C.E.C.) in milliequivalu-
of the six soils tested were found to increase by ti' •
also found increasing sorption rates to be correiai <
temperature, wind velocity or exposure time.
Hanawalt (1969b) evaluated the influence of st
such as organic matter, pH, cation exchange capacit •-•
moisture on NH sorption rates. The property found •
sorption rates was the soil moisture content, betwi.••••'
a positive correlation. Hanawalt concluded then in•*
Nil, is absorbed by a soil is affected most "by thu
-------
-82-
the rate of supply of anmonia to the soil surface," namely the atmospheric
factors mentioned above, and "the fraction of the ammonia molecules which
are actually captured," which is dependent upon the moisture content
and air permeability of the soil.
Water Bodies: Because the absorption of atmospheric ammonia by lakes
and streams can promote eutrophication in that water body, a great deal of
concern has developed, especially in areas adjacent to large cattle
feedlots and sewage treatment plants which both release great quantities
of ammonia into the atmosphere. Hutchinson and Viets (1969) showed that
the ammonia volatilized from cattle feedlots was absorbed by nearby
water bodies. In their study, they calculated that Seeley Lake in
northeastern Colorado, which is about 2 km from a 90,000 unit feedlot,
absorbs enough Nil from the air over a one year period to elevate its
nitrogen concentration by 0.6 mg/liter. An interesting result of their
study was the discovery that although there was a tendency for the NH
content in precipitation to increase as the feedlots were approached,
in no instance at any site, near or far from a feedlot, was the NH
content in precipitation very high. It appears that in all cases the
amount of NH removed from the atmosphere by precipitation was "insignifi-
cant compared to the amount absorbed directly from the air by aqueous
surfaces in the vicinity of cattle feedlots" (see Table XIV).
In an attempt to quantify the removal of NH released from a point
source (such as from a plume produced by a sewage treatment plant) by an
aqueous surface, Calder (1972) developed a simple mathematical model
which takes into consideration the atmospheric transport and diffusion
o£ the plume as well as a characterization of the removal process,
-------
-83-
Table XIV. NH* - N ABSORBED BY SURFACE WATERS AND ITU
ADJACENT TO A CATTLE FEEDLOT
* Mean weekly absorption of ammonia-nitrogen by di!
the period 27 July 1968 through 27 February 1969
where measurement began 27 September 1968)
t Estimated annual absorption of NH - N by water ^
dividing mean weekly rates by two and multiplying
are divided by 2 because that is the ratio of NH^
acid traps to lake water traps)
§ NFL - N in precipitation measured during the perj-.
21 November 1968, Total amount of p'pt. over thj:
site 1 and between 3,0 and 3.6 cm at the other sii
of NH* measured in the precipitation was extrapoJ,
to give the measurements listed.
Site Site description
1 No feedlots or irrigated fields within i> ' " •' .:
3 km; no large feedlots or cities within 15 km
2 Only small (less than 200-unit) feedlots
within 4 km, none closer than 0.8 km
3 About 0.2 km east of 800-unit feedlot and
0,6 km west northwest of another similar
feedlot
4 On northeast shore of Clark Lake and 0.5 km
southwest of 9,000-uait feedlot
5 On southeast shore of Seeley Lake and 2 km
west northwest of 90,000-unit feedlot
6 About 2 km east of 90,000-unit feedlot ; '• to
7 About 0,4 km west of 90,000-unit feedlot ' -I
-------
-84-
in this case the deposition velocity. The reader is referred to his
article for the mathematical development of this model. Given the
uncertainty in the value of the deposition velocity for NH, absorption
by water, Gaidar concludes that absorption of NH from an NH, cloud,
»J *5
passing over a water body some 30 km long at an average speed of 5 m/sec,
could exceed 20 per cent.
Rainout and Washout: Ammonia, as previously mentioned, is extremely
soluble in water. Once dissolved it ionizes to NH., the amount depending
on the pH, as follows:
. NH3'-H20 £ NH* + OH"
The rate constant equals J..774 x 10 sec (Robinson and Stokes, 1959).
Normally the pH of cloud water ranges from 5 to 6. Under atmospheric
conditions, assuming that an equilibrium solution is attained and CO is
present in normal concentrations, Junge (1963) has shown that approxi-
mately 87% of the NH, present in air at a concentration of 3 yg/m could
be absorbed in cloud droplets. Unfortunately, equilibrium is not usually,
if ever, reached in cloud droplets.
Ammonia has been shown to be an important catalyst for the oxidation
of SO and NO in solution (van den Heuval and Mason, 1963). Both Scott
and Hobbs (1967] and Miller and de Pena (1972) have shown that as the
partial pressure of NH in the atmosphere increases, greater concentra-
tions of SO can be dissolved and oxidized in solution to form sulfate
particles. The resulting aerosols, if evaporated, are composed largely
of (NH ) SO particles with probably minor amounts of NH HSO,, NH HSO.
and (NH4)2SO (Miller and de Pena, 1972). In the case where NH is
-------
co-absorbed with NO , the resulting particles would be composed
primarily ol NH NO .
It has been estimated (Robinson and Robbins, 1968) that almost 7S'i
of atmospheric ammonia is removed from the atmosphere by com-eision to
NH, lonb which condense in cloud droplets or particles and may e\rapora('
to form aerosols, Georgii (1963) found that in a polluted a.ir mass
containing approximately 500 ugSCL/m and 10 ugNH_/m , 98% of the anunon '
is converted to (NH.)?SO. aerosols. McConnell (1973) has estimated tluj'
of a total NH source strength of about 17,4 x 10 kg NFL, N/yr,
*J -_)
approximately 3 x 10 kg of this ammonia is removed each yeui by
rainout. He points out that this is a very conservative estimate and
"may be low by as much as a factor of ,3." Earlier, Robinson and Robbif
estimated that 280 x 10 kg NH,, - N are removed each year by precipita
O
tion. On the whole, quantitative estimates are rather sparse and show
4
a great discrepancy. To amend this, extensive measurements of the Nil
content of rainwater must be made and analyzed,,
Atmospheric Reactions Involving NH : Until just recently the re^
involved in the destruction of atmospheric ammonia were poorly define-i
McConnell (1973) has suggested a number of reactions he believes to br
significant in its destruction.
In the stratosphere the primary destructive reaction involves the
photodissociation of NH by sunlight with wavelengths greater than
O
2300 A. The basic reaction is
hv -* NH + H
-------
-86-
-4 -1
The reaction rate is 1.3 x 10 sec . The various excited NH2 and NH
states possible, together with their observed threshold wavelengths and
their respective quantum yields are listed in McConnell's Table 1.
Great quantities of hydroxyl radicals are now believed to exist
in the troposphere (Levy, 197]), the number density appearing to be on
12 -3
the order of about 10 m (McConnell et al,, 1971). Reaction of NH
with these OH radicals in the troposphere and lower stratosphere is
therefore possible, and, probably constitutes an important sink as well.
The reaction is as follows:
NH + OH ->• NH + HO
-19 3 -1
Stuhl (1973) records the rate as 1.5 x 10 m sec at room temperature.
The magnitude of this rate constant is consistent with observations of
Worley, Coltharp and Potter (1972) and Albers, Hoyermann, Wagner and
Wolfrum (1969). McConnell points out that this same reaction can and
probably does occur in the stratosphere, but it is not as effective
there as photolysis. He further points out that although reactions of
NH, with 0, 0( D) and 0 are possible, evidence suggests that such
reactions would not play a significant role in the destruction of ammonia.
According to McConnell, "the most likely atmospheric products"
of NH destruction are N_ and NO. In his paper he lists those reactions
most likely to occur along with any information that might be available
on their reaction rates. An anomaly exists here that cannot as yet
be answered. As McConnell indicated, oxidation of NH_ can be the source
*!)
of a significant quantity of nitrogen oxides, especially NO (estimated
at 7.04 x 10 kg N0y per year as N). But, he points out that the
-------
-87-
magnitude of this NO source can be reduced due to a possible immediate
reaction of NO with NhL producing N and HO. This then would mean that
oxidation of NH could constitute a sink for NO. In effect then, oxida-
o
ti'on of NH can both produce and consume NO. The proportion of NO
produced to that consumed by these reactions if they do occur is not
as yet clear, but his contention is that the NO produced far outweighs
that consumed.
Environmental Ammonia Cycle
Relatively little quantitative information is available on the
strengths of both sources and sinks of ammonia. In fact, research into
possible atmospheric oxidation reactions is so recent that any estimates
made are subject to gieat uncertainty. Separate atmospheric ammonia cycles
have never been constructed; its cycle is always considered together with
that of nitrogen oxides. Because not very much is even known about the
nitrogen oxides cycle, the uncertainty involved in the combined cycle
must be very high.
To date the only nitrogen compound cycles devised are those of
Robinson and Robbins (1968, 1970) and McConnell (1973) and a total
geochemical nitrogen cycle of the present research, The first three
cycles are compared in Table XV. The discrepancy among the three in
source and sink estimates is immediately obvious. More research must
be undertaken to provide us with a better understanding of the part
ammonia plays in atmospheric chemistry and to quantify possible sources
and sinks such as the oceans, lakes and vegetation.
-------
-88-
Table XV. ATMOSPHERIC FLUXES INVOLVED IN VARIOUS AMMONIA CYCLES
(in 1010 kg N/yr)
Robinson Robinson
§ Robbins* § Robbins
(1968) (1970)
I. Sources
a. anthropogenic .35
b. biological 670± 95.7
TOTAL 670 96.05
II. Sinks
a. precipitation 280 18.6
b. dry deposition 70 4.9
c. oxidation to NOY (troposphere) N.E.
A
d. oxidation § photolysis to NO N.E.
(stratosphere)
e, gaseous deposition 90 74.9
TOTAL 440 98.4
McConnell
(1973)
0.4
17
17.4
3
7
7
0.04
17.04
N.E. = Mechanism recognized but no estimate made.
*NOTE: (see note in Table XI)
ISource strength here was adjusted by Robinson and Robbins to
provide an additional amount of nitrogen needed to balance other
portions of their nitrogen compound cycle.
-------
-89-
GLOBAL NITROGEN CYCLE
The role which nitrogen compounds play in the chemistry of the
atmosphere and earth is not well understood. A summary of what is known
and has been suggested about the routes and rates of transport of these
compounds is shown in Figure 4,
The diagram has been divided into certain spheres or reservoirs in
which there exists an equilibrium concentration, or burden, of nitrogen.
These burdens have recently been compiled and, in some cases, calculated
by Wlotzka (1972). The only elaboration made upon his tabulations con-
cerns the abundance of the nitrite and ammonium ion present in the ocean.
Sverdrup, Johnson and Fleming (1942), as cited by Wlotzka, calculated the
NO" - N content of the ocean to be 5.7 x 10 Tg (5.7 x 10 7 g). Emery,
O
Orr and Rittenberg (1955)hkve estimated the total inorganic combined
nitrogen content, that is, N0~ N0~ and NH* to be 5.8 x 10° Tg. Therefore
in this cycle the N02 and NH inventory of the hydrosphere has been set
equal to 0.1 x 10 Tg. These burdens, shown encircled in their
12
respective reservoirs in Figure 4, are given in units of Tg N (10 g N).
There is so little quantitative data available concerning rates of
transfer, or fluxes, of these compounds between reservoirs it is impossible
to even attempt to balance this cycle. As a result the cycle suggests
the possible routes of transfer and offers those fluxes which have been
12
estimated. Fluxes are expressed in Tg N/yr (10 g N/yr) and the direction
of transport is indicated by arrows.
Details of the atmospheric portion of this cycle, with the exception
of biological nitrogen fixation and denitrification as a source and sink
of N respectively, and the volcanic emanations of N , have been discussed
-------
-90-
2 z NOUdaOSOV
y oooi - ooi
Nouwxb 'oia
u
X
u
c;
cd
o
1—J
CD
E-
«^-
0)
tifl
•H
tLi
-------
cocunc removed CI*«'%T \,h; .< < •
bunion :/. NO., and ' r.1 ^ ':... ( ''-
and this is noted on the dia^i-jp.
In formulating a rutrogen bucifi. ; •; L i. , . • ,• i
aiso estimated that 8 Tg N/yr j.- at' '-. •• • >< • .t :
fixation, probably via some species of b im -^ieui aiv
An additional contribution to the t.-.-.,^ n -; ..•, , - ,
is surface and subsurface 2 and dru • j.if',!.
19 Tg N/yr are added to the ocean .,}< :,u-".' '..;
(1972) has caj.culatea this co<'ur.i:), • u. ... i
The most widely recognised prci « . ,.i,. ;
the ocean is sedimentation The <-'-r- ' r,, . L •
manner has been estimated to range • i". • ,• ,/, /> •
8,6 Tg/yr (Emery et al , 1955) 0. h- , -•:• •,...
include denitrification, producing .\ 'i-J \,'.'
-------
-92-
Liss and Slater (1974) calculated the flux of NLO from the sea to the air
at 38 ig/y.- (120 Tg N_0/>r). Wlotzka estimated this flux to range between
20 and 130 Tg/yr. The amount of N~ released from the ocean as a result
of denitrification is unknown, as is the amount of NH_ which escapes.
Nothing is known about the contributions of mantle degassing,
volcanism, and weathering of the lithosphere to the total amount of
nitrogen carried to the ocean by land drainage. It has been estimated
though that about 0.1 Tg/yr is released to the atmosphere by volcanic
activity.
A great deal of nitrogen containing fertilizers are applied each
year to the soils across "he world. Wlotzka recently estimated it to
total 30 Tg/yr. Not only does this nitrogen play an important role in
the nutrient cycle between soil and plants, but much of it is subject
to leaching and would therefore contribute to the amount of nitrogen
removed by surface and subsurface land drainage, ultimately to wind up
in the sea,
As mentioned previously, the quantity of nitrogen present in the
ocean is constant with tiir.e. Based on the nitrogen budget of sea
water as formulated by Emery et al . (1955), Dugdale (1972) calculated
the residence time of nitrogen in the ocean to be:
5
86 Tg/yr
,.,0 9.2 x 10 Tg In4
T = M'R = * g 10
or about 10,000 years which is an extremely short period of time on the
geological time scale. For the sake of comparison, it was earlier
calculated that the residence time of sulfate in the ocean is on the
order of millions of years.
-------
-93-
Figure 4 and the explanation above have attempted to show how vcix
little is known about the global nitrogen cycle. Each question mark ii
the diagram needs to be replaced by a number representing an actual I i
in the environment. Those fluxes that have been estimated need confine
tion. Until all that is accomplished, any nitrogen cycle devised can -
be, at best, conjectural.
-------
-94-
OXIDANTS
OZONE:
The problems involved when significant amounts of ozone are present
in the atmosphere have come to light probably more as a result of the
photochemical pollution problem in Los Angeles than from any other single
factor. While background concentrations of ozone probably range from
about 20-60 yg/m (0.01-0.03 ppm), in urban centers like Los Angeles,
it is not unusual to have 0 present at levels greater than 500 yg/m
O
(.25 ppm). Ozone levels up to 400 yg/m (0,2 ppm) usually will not cause
any deleterious effects (Masters, 1971), but at concentrations of 600 yg/m
(0=3 ppm) ozone causes irritation of the mucous membranes in the nose
and throat. At. somewha- higher levels, it can cause coughing, choking
and severe fatigue. When present at relatively high levels, such as
those that occur in severe photochemical smogs, ozone causes bronchial
irritation and interferes with normal lung functioning, causing breathing
difficulty and chest paini. The highest ozone concentration detected in
the Los Angeles atmosphere was 2,000 yg/m (0.99 ppm) in 1956 (Chambers,
as cited in Tebbens, 1968).
Another serious problem related to the occurrence of atmospheric
ozone is its toxic effect on vegetation, especially field and forage
crops (such as tobacco), ^.eafy vegetables, shrubs, fruit and forest trees
(particularly conifers).
Sources
The maximum ozone density occurs at an altitude of 25(±5) km, in
the mid-stratosphere. This ozone though is not formed there but between
30 and 60 km where the following reaction proceeds:
-------
-95-
0+0 + M + O+M
£. O
In this altitude range though ozone is relatively unstable. It may be
destroyed either by collision with atomic oxygen to form molecular oxygen
or by photolysis. The cycle is continuous; formation, destruction,
formation. . , As a result an approximate state of equilibrium exists
above about 40 km. The accumulation of ozone at 25 km is a result of i'<
downward transport to a location where its destruction is less likely
(Barry and Chorley, 1970).
Because the wavelengths of ultraviolet radiation that penetrates
the troposphere are too long to cause photodissociation of oxygen, it
has long been accepted that the presence of ozone in the troposphere is
due primarily to the transport of ozone down from the stratosphere. Th ;
amount present in the troposphere would then be related to the injectun
rate through the tropopause, estimated as ranging between 1.9 and 7.5
x 10"5 kg/yr (Junge, 1962).
Recently, Chameides and Walker (1973) proposed a model wherein both
seasonal and diurnal variations in the tropospheric ozone density are
assumed to be caused by photochemical changes rather than a change in
the flux of stratospheric ozone-rich air into the troposphere.
Their model calls for the production of 0 by the methane oxidation
J
scheme suggested by Crutzen (1973) . The oxidation of methane produces
hydroperoxyl radicals which then react with nitric oxide:
H02 + NO ->• OH + N02 followed by (} >
NO + hv -* NO + 0 and (/ i
-------
-96-
0 + 02 + M + 03 + M C3)
At 27°C, k is about 5 x 10" m sec" (Levy, 1973a). Based on calculated
atmospheric densities of HO and NO , Chameides and Walker calculated an
12 -3 -1
ozone production rate of approximately 5 x 10 m sec , and photo-
chemical lifetimes of a few tenths of a day near the ground, of one day
at 5 km, and 10 days at 10 km. They concluded that ozone is in a state
of photochemical equilibrium in the troposphere.
Upon developing a photochemical model of reactions affecting the
density of tropospheric ozone, they found they could satisfactorily
reproduce the seasonal variations found in ambient ozone densities (see
Figure 5) by applying their model to temperature and relative humidity
conditions typical of winter and summer. By analyzing the specific
effects that lowering the temperature, humidity and photodissociation
rates each make in altering the summer profile to one characteristic of
the winter, they conclude that decreasing the photodissociation rates
alone causes a realistic alteration of the summer profile to that of
winter (Figure 6). This model shows that photochemical processes rather
than the vertical transport of 0 -rich air can be responsible for the
o
presence of ozone in the troposphere. It also presents a challenge for
other researchers to prove whether or not most tropospheric ozone
originates in the stratosphere as commonly believed or in the troposphere
by the reactions these authors propose.
Removal Mechanisms
Ozone is relatively insoluble in water; at 20°C and 1 atm its solubi-
lity is 0.052 g/lOOg HO. Therefore, removal of ozone from the atmosphere
-------
-97-
^ ^^
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-------
-98-
by washout and rainout can be disregarded. There is evidence though that
ozone is removed by the oceans. Ozone is also absorbed by vegetation and
soil and these appear to represent major sinks, for 0 . Due to its
o
nature as a strong oxidising agent, ozone participates in a number of
atmospheric reactions, especially in polluted atmospheres. Junge (1962)
has estimated the tropospheric residence time of ozone to range between
3 and 4 months.
Vegetation: Ozone has caused a great deal of damage to vegetation
in the United States and its effects on both plants and the economy have
been well documented (Rich, 1964 and Millecan, 1971). When ozone is
present in the atmosphere at low concentrations, vegetation is capable of
absorbing some without inflicting damage upon itself. Hill (1971) showed
that an alfalfa canopy removed ozone from his experimental chamber at a
rate of 1.7 cm/sec. In this study the threshold level above which the
-43 3
plants sustained injury was 4.16 x 10 mol/m (200 yg/m or 0.1 ppm).
Previously, Hill and Littlefield (1969) had shown that high concentrations
of ozone can cause stomatal closure in the leaf thereby reducing the
uptake of ozone. Unfortunately in Hill's study the injury threshold limit
-4 3
of alfalfa was 4.16 x 10 mol/m which was also the level above which the
stomates began to close. As a result, the effect of stomatal closure on
ozone uptake rates could not be adequately evaluated.
The hypothesis that the uptake of 0 is dependent on the degree of
O
stomatal opening received additional support by Rich, Waggoner and
Tom1inson (1970) as a result of their experiments on bean plants. They
concluded that ozone follows the same but reverse path as that of water
vapor which is transpired by the leaf. The water vapor originates from
-------
-99-
saturated cells directly beneath the stomata and is released when the
stomata are open. Conversely, the ozone travels through the open
stomata onto the surface of the substomatal cells where it is reduced to
very low concentrations. Hill points out that although ozone's solubility
in water is not very high, it breaks down relatively rapidly and therefore
a fairly high uptake rate is to be expected.
Soil and Water Bodies: The ground has long been recognized as a
sink for ozone. Over the last two decades several estimates of the
vertical flux of 0 into the ground have been made. Junge (1962) tabulated
the estimates made through 1962 [see his Table I]. For the most part,
-12 2
the estimates range from about 3 to 200 x 10 kg 0 /m -sec. Jurige calcu-
-12 2
lated the flux density to be 60 x 10 kg/m *sec which is in good agreement
-12 2
with the estimate of 48 x 10 kg/m 'sec made by Kroening and Ney (1962).
These measurements of 0 into the ground include destruction by both soil
and vegetation. No attempts were made to differentiate the ozone destroyed
by each, and no suggestions were offered as to what the exact reason for
the destruction at these surfaces were.
Aldaz (1969) was the first to offer data on the ozone flux into various
surfaces: soil, vegetation, snow and water (Figure 7). These fluxes were
determined assuming a partial ozone density of 40 yg/m (0,02 ppm, or
natural background concentration levels). Using an equation analogous
to that used by Liss and Slater (1974) to estimate the flux of a gas
across the air-sea interface, Aldaz calculated the flux of 0 into both
the terrestrial and marine portions of the northern and southern
hemispheres. The equation used to determine the ozone flux is
F = 1.25 x 1010 kq
-------
-100-
10
16
o
0>
CM
l|015
to
o
O
X
o>
c
o
5
10'
10
13
^JUNIPER BUSH (N.M.)
^-TUNDRA(ALASKA) KELLEY, 1968.
— SAND OR DRY GRASS (N.M.)
^- GRASS (AUSTRALIA) GALBALLY, 1968.
— GRASS (NEBRASKA) REGENER, 1957.
— SNOW (N.M.)
WATER(FRESH)(N.M.)
WATER(SEA)(40°N, 70°W)
WATER (SEAKTROPICAL, MIDDLE LAT.)
1VATER (DISTILLED)
CLEAN MYLAR
Figure 7. Ozone fluxes; into different surfaces assuming a partial
ozone density of 40 yg/m^.
-------
-101-
where q is the partial density of 0 in yg/m , and k, the j" • i
is a measure of the chemical reactivity of the surface to -^i',"
then, k is actually a deposition velocity term. The react !••>
used were 0.60 cm/sec over land, 0.04 cm/sec over water aii" '
snow. Data Aldaz used to compute the global ozone sink i s :.- •
XVI. Because there are no measurements of ozone uptake by > • i
tion, Aldaz assumes in estimate (A) that the rate of destriu . •
tropics is equal to that in other land regions, and in (BJ ;,
destruction rate 5 times greater in the tropics than in ot!:<
In conclusion, he found the sink strength of the earth's si',1 -•
between 1.3 and 2.1 x 10 kg/yr. He also showed that de';l<'
ozone by New Mexican soils is about 15 times faster than ». -, •
Ocean (Figure 7), and that bare, dry soil destroyed approxi'-1
more 0 than when moist It must be pointed out that the =, -. :
which the soil destroys ozone was not identified.
Turner, Rich and Waggoner (1975) recently found thcil • .•
freshly cultivated, fine sand loam removed ozone at a iat._ -'••:
-8 2
.5 - 2 x 10 mol/m -sec when present in the atmosphere uf
ranging from 64-208 ng/m . They also found that as the soi i :,.
content was increased, the resistance of the soil to 0,, !••"• i . i
increased. As mentioned above, Aldaz found a similar dir.i...
in the sand loam soil he tested in New Mexico; i.e., as li<
increases, ozone destruction by the soil decreases,
Atmospheric Reactions Involving 0 : Although react.-.
with natural contaminants such as terpene do occur
1965; and Ripperton et al . , 1967), its importance arises I
-------
-102-
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-103-
*
photochemical reactions in which it participates in poll-"'
namely the photo-oxidation of hydrocarbons in the preben ,
dioxide. The initial oxidation of olefins by ozone foi *<•
to a long series of reactions which produce ketones, al-f1'
acids and nitrogen-containing compounds such as peroxy;s ••>
The presence of these compounds in the atmosphere has !•• -."•
cause of considerable eye irritation (Schuck and Doyle, i
constants for the reaction of ozone with numerous hydr-< ••
summarized by Altshuller and Bufalini (1971) and BufaJ < '
(1965). Because the reactions of ozone in polluted a1 ir <
numerous and have been the subject of extensive invest i;; '•
review, it would not serve the best interest of this p:: ••
reactions here. For additional information the readei
Hecht and Seinfeld (1973), Dutsch (1971), Ripperton an,! ,
Stepnens (1969), Altshuller and Bufalini (1965), and I... • •
-------
-104-
ORGANIC GASES
Organic gases constitute the second major group of air contaminants.
This group includes all classes of hydrocarbons and those formed when some
of the hydrogen of the original compound is replaced by other substituent
groups containing nitrogen, sulfur, or oxygen. Organic gases are further
subdivided into reactive and non-reactive classes. Among these classes
only hydrocarbons are cor side red here. The more important reactive
hydrocarbons include the olefins and aromatics. Paraffinic hydrocarbons
are classified as non-reactive,
REACTIVE HYDROCARBONS
Because of their role in photochemical reactions in polluted
atmospheres, reactive hydrocarbons have been the subject of great interest.
These photochemical reactions produce smog which is associated with eye
and respiratory tract irritation, reduced atmospheric visibility, and
plant damage.
Limited ambient data from Point Barrow, Alaska (Robinson and Robbins,
1968)indicate that ethylene, the most abundant hydrocarbon of this group,
is present at a concentration of less than 1 ug/m (less than 1 ppb).
In the absence of other ambient measurements one might consider the
concentration of ethylene measured at the above location to represent
the upper limit of the background concentration for components in this
group„
Sources
A variety of hydrocarbons are released to the atmosphere as a result
of both anthropogenic activity and natural processes. The most important
-------
-105-
anthropogenic source of hydrocarbons is motor vehicle exit,:.'
incomplete combustion of fuel. Robinson and Robbins (llJbSi
that the total annual emission of olefins and aromatics tV;
9
tion of various fuels is 27 x 10 kilograms. A bibliograj.i
emission sources has been compiled by the U.S. Department <•>
Education and Welfare (1970).
Plant species also release appreciable quantities or
organic substances to the surrounding air. The major rc.i • •
emitted by trees are ethylene, monoterpene (C,n), and isop,
a recent study, Rasmussen (1972) concluded that the
9
a global natural source of 175 x 10 kg of reactive hydiv*
year. This emission rate is 6 times greater than that csi . •
reactive hydrocarbons of anthropogenic origin,
Removal Mechanisms
Hydrocarbons in general are not water soluble, and ' !
cannot be directly removed from the atmosphere by wet pix» >
washout and absorption by surface waters, Various studios '
that photochemical reactions are important in removing roa'
carbons, although the products formed may cause detriment.1 '
such as eye and throat irritation. This class of hydroca.i'
their emission into the atmosphere undergo rapid chemical >'
in the presence of: atomic oxygen and ozone (Bufalini and '
1965); oxides of nitrogen (Schuck, 1961; Alley, Martin and
ozone and sulfur dioxide (Cox and Penkett, 1971b) and; n.i i
and sulfur dioxide (Schuck and Doyle, 1959)
-------
-106-
The basic kinetic mechanisms of hydrocarbon reactions in the
atmosphere are given by Hecht and Seinfeld (1972). These authors present
a 15-step mechanism for photochemical smog formation, with rate constants
and stoichiometric coefficients chosen according to the particular hydro-
carbons involved in the reactions and the initial reactant ratios. The
state of the art of photochemical reactions is analyzed by Dodge (1973)
and Seinfeld, Hecht and Roi:h (1973). The reader is referred to these
publications for an up-to-date view.
Quantitative data on the rates of these atmospheric reactions are
rare. The limited data available as reported by Hidy (1973) suggests
that 1-10% by weight of the; reactive hydrocarbons emitted into the
atmosphere are converted to aerosols and eventually removed by scavenging
or deposition. The remaining hydrocarbons are eventually oxidized to
carbon dioxide and water vapor.
Smith et al. (1973) investigated the capacity of soils to absorb
ethylene and acetylene. Sterilized and air-dried unsterilized soils
removed neither of these hydrocarbons from the test chamber. Moist
unsterilized soils (50% saturated) did reduce the ambient concentrations
of these gases although the uptake rate was relatively slow compared
to the other gases they studied (S0_ and CO). They concluded, as did
Abeles et al. (1971) in an experiment on ethylene uptake by soil, that
the sorption of both ethylene and acetylene is due to microbial activity
in the soil. Smith et al. found that the soils they tested removed
-9
ethylene at average rates ranging from .14-.97 x 10 mol per gram of
-9
soil per day (mol/g*d) and acetylene from .24-3.12 x 10 mol/g*d.
-------
-107-
NON-RL ACT[VL IIYDROCARBONS
This group, which consists of methane and the higher saturated
hydrocarbons, has been found to be much less involved in photocliunii1,, i
reactions and smog formation than reactive hydrocarbons. By jar ! lir •
abundant paraffinic hydrocarbon in the atmosphere is methane, i'letb;-h:
background concentration is about 1000 yg/m (1.5 ppm) , while Lh,/
background concentration for heavier gases in this class is ie: ,s I h,n)
1 yg/m" (less than 1 ppb) (Robinson and Robbins, 1968),
Sources
The anthropogenic source of paraffinic hydrocarbons is the in-i.,
combustion of fuel in motor vehicles. The annual emission o! i-,ira!i i<
9
hydrocarbons is estimated to be 60 x 10 kg (Robinson and Kolihui:, i
9
Among natural sources approximately 310 x 10 kg of medi-iiie i.-.
produced annually in swamps and various water bodies as a revtlt «i
bacterial decomposition. The relatively high concentration of moth.,'.
in the atmosphere compared to other organic gases is related to tin--.
natural process.
Removal Mechanisms
Because they are so insoluble in water, paraffinic hydroea i }jf>.<
can not be removed from the atmosphere by wet processes. The J.H'IJVJ; ,
sink for methane in the troposphere is its oxidation by hydroxyl i;),;:•
to form carbon monoxide. The initial reaction is as follows:
CH + OH -> CH + HO
• O *-
-------
-108-
-12
The rate coefficient for this equation equals 5.5 x 10 exp (-1900/T)
(Grciner, i970) . For a complete development of this oxidation scheme the
reader is referred to Levy (1971), McConnell et al. (1971) and Levy
(1972, 1973a). Based upon density profiles of CH and hydroxyl radicals
in the troposphere, and the rate equation given above, Levy (1973b)
calculated the average daily loss of methane at a particular altitude.
_ Q
The total column loss rale for methane was found to be 7.48 x 10
2
mol/m -sec. This results in a tropospheric residence time for methane
of 2 years.
Rasmussen, Hutton and Garner (1968) and Robinson and Robbins (1968)
have also suggested that the volatile organic components of the atmosphere
are removed by bacteriological processes and vegetation. However,
quantitative data and rate equations for these removal mechanisms are
nonexistant at present.
-------
-109-
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TECHNICAL REPORT DATA
>'L jw r':t;JI>-i"jj-,tn on the rc;'t'rse before complsti'ig)
J3, RECIPIENT'S ACCESSION-NO.
u; ;Hd".;,,r.'/. Removal Processes for
p'ieric rollutants
Kabel
Mamoo?1 Faheri
(Pahlavi Univ., IRAN)
*v:,NG OH "ANIMATION NAMb AND ADDRI-.JS
•nr for Air Environment Studies
renske Laboratory
Pennsylvania State Uni versitv
•ersity °ark. PA 16802
J5. REPORT DATE
j_
6 PE.RF-ORMING ORGANIZATION CODE
8 PfcHI-OHMING ORGANIZATION HI COM I MO
CAES Publication No. 367-74
•• TOfiiNG AGENCY NAMfc AND ADDRbSS
;eorology Laboratory, EPA
/onal Environmental Research Center
ectrch Triangle Park, N. C. 27711
10 PPCiGRAM ELEMENT NO.
1A1009
11. CONTRACT/GRANT NO.
800397
K3. TYfp OF REPORT AND PERIOD COVERED
I terim
j 14. SI?ONS^OR.TN"G AG¥^Jcv CODE
MOTES
••I. A3STRACT
T.his review attempts to briefly "llustrate what the "state of the art" is in the
recognition of the various soum;s and natural sinks of gaseous pollutants. The
removal mechanisms include absorption by vegetation, soil, rock and water bodies,
precipitation scavenging, and chemical reactions within the atmosphere. The
nature and magnitude of anthropogenic and natural emissions of the gases
considered i%S, SCU NO, NO, NO , NH0, CO, 0 , and hydrocarbons), along with
r,revr ambient^backgrouna concentrations' and information on their major sinks
•identified to date, are discussed. In the case of sulfurous and nitrogenous
cor.iooundSj this information r-.as neen used to prepared total geochemical cycles.
DESCRIPTORS
KEY WORDS AISiD DOCUMENT ANALYSIS
jb. IDENTIFIERS/OPEN ENDED TERMS
Hydrogen sulfide
Sulfur dioxide
;.". reactions Carbon monoide
• ;".cr= NitrotiS oxide
Nitric oxide
Nitrogen dioxide
Ammonia, Hydrocarbon
Sinks
Sources
Natural removal
i: CVClSS
:yc! ~
uzone
O.\ STATEMENT
nn ted
,19. SECURITY CLASS (This Report)
; Unclassified
J20. SECURITY CLASS (This page)
\ Unclassified
c. COSATi 1'iclci/Group
21. NO OF PAGES
22
122
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