Ecological Research Series
Chemical/Biological  Relationships
Relevant to Ecological Effects
of  Acid Rainfall
                                National Environmental Research Center
                                 Office of Research and Development
                                 U.S. Environmental Protection Agency
                                       Corvallis, Oregon 97330

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                               JUNE 1975
               John 0. Reuss
  National  Ecological Research Laboratory
  National  Environmental  Research Center
         Corvallis, Oregon  97330
        Program Element No.  1AA006
           ROAP/Task No. 21ALU32
         CORVALLIS, OREGON  97330
    For sale by the Superintendent of Documents, U.S. Government
          Printing Office, Washington, D.C. 20402


     This paper deals with problems  concerning measurements  of rainfall
acidity and interpretation in  terms  of possible  effects  on the soil-
plant system.

     The theory of  acidity relationships  of the  carbon dioxide-
bicarbonate equilibria  and its effect on  rainfall  acidity measurements
is given.  The relationship  of a  cation-anion balance model  of acidity
in rainfall to plant  nutrient  uptake processes  is  discussed, along with
the relationship  of this  model to a  rainfall  acidity model previously
proposed in the literature.  These considerations  lead to the conclusion
that average  H  concentration  calculated  from pH measurements is not a
satisfactory  method of  determining H  loading from rainfall  if the rain
is not  consistently acid. Calculating  loading  from H  minus HCO-",
strong  acid anions  minus  basic cations,  or net  titratable acidity is

     The flux of  H+ ions  in  soil  systems  due to  plant uptake processes
and sulfur and nitrogen cycling is considered.   H   is produced by
oxidation of  reduced  sulfur  and nitrogen  compounds mineralized during
decomposition of  organic  matter.   Plant  uptake  processes may result in
production of either  H+ or OH" ions.  Fluxes of  H   from  these processes
are much greater  than rainfall H+ inputs, complicating measurement and
interpretation of rainfall effects.   The  soil acidifying potential due
to the  oxidation  of the MM/ in rainfall  is examined, with the conclusion
                          4                                        +
that acidity  from this  source  is  of  a similar magnitude  to direct H
inputs  common in  rainfall.

     This report  was  submitted by the National  Ecological Research
Laboratory under  the  sponsorship  of  the  Environmental Protection Agency.
Work was completed  as of  January,  1975.

Sections                                             Page
     I   Conclusions                                   1

    II   Recommendations                               3

   III   Introduction                                  4

    IV   Carbon dioxide - bicarbonate equilibrium
         in rainfall                                   5

     V   Cation-anion balance in relation to acid
         rainfall                                     18

    VI   Acidity relationships of the sulfur cycle    26

   VII   Acid-base relationship of the nitrogen
         cycle                                        34

  VIII   References                                   44

                      LIST OF FIGURES


The relationship between pH and HC03~ activity in
equilibrium with atmospheric C02  (316 ppm).                  9

Plot of the pH and HC03~ levels reported by
Egner and Eriksson (1955b, data set D501.).                 13

Diagram showing the effects of varying C02
partial pressure on pH.  Dotted line indicates
316 ppm C02 and numbers refer to  pH at 316 ppm              15

Relationships between net acidity as calculated
by anions (excluding HCOo~) minus basic cations,
and observed H  minus HC03  for data set reported
by Pearson and Fisher (1971).  Coastal stations
were not included.                                          21

Relationship between net acidity  as calculated
by anions (excluding HCCL"") minus basic cations,
              +         J _
and observed H  minus HCCL  for data set D502
reported by Egner and Eriksson (1955b).  One
coastal station not included.                               23

Simplified sulfur cycle showing acid formed or
consumed by the various processes.                          28

Simplified nitrogen cycle assuming all N utilized
in the NHL  form, and showing acid formed or
consumed by the various processes.                          35

                               SECTION I


     A number of chemical and biological considerations relevant to
understanding the effects of rainfall acidity on soil-plant systems have
been discussed.  The following points are particularly important and
should be thoroughly understood by researchers engaged in acid rainfall

     1.   Hydrogen ion inputs from rainfall to soil and aquatic systems
          calculated from "weighted average" pH measurements without
          regard to bicarbonate levels or carbon dioxide partial
          pressures may be seriously in error.  Alternative methods
          include; H  minus HCO%> net titratable acidity, and acid
          anions minus basic cations.

     2.   The potential of rainfall to acidify soil or waters is
          characterized by its containing more acid anions than basic
          cations.  This excess of anions can be an appropriate measure
          of the capacity of the rainfall to contribute to the acidity
          of soil-water systems.

     3.   The acidity of a plant growth medium is affected by the
          relative uptake of cations and anions by the plants growing in
          the medium.  The overall effect of acid rainfall is not only
          the effect of H+ ions added to the system but may also reflect
          the relative plant uptake of the various ions added to the

4.   In the natural sulfur cycle, H+ ions are continually being
     formed and consumed by various biological processes.  An
     important aspect of the acid rainfall question involves the
     effect of additional sulfur loading on these processes.

5.   Soil nitrogen transformations normally involve a substantial
     production and consumption of H  ions.  This flux of H  ions
     is much larger than that involved in sulfur transformation or
     the input from acid rainfall.  Measurement of the effects of
     H  inputs from outside the system may be complicated by this
     larger flux due to  normal nitrogen transformations.
     The oxidation of ammonium to nitrate by aerobic chemauto-
     trophic bacteria results in the production of acid.  Oxidation
     of the ammonium contained in rainfall may result in an H
     production of a similar order of magnitude to the direct input
     of H  in acid rain.  The consideration of this process sub-
     stantially complicates the interpretation of rainfall induced
     H  loading.

                              SECTION II


1.   Hydrogen ion concentrations (pH) measurements on rainfall should
     normally be made after equilibration with known COp partial pressures,
     preferably the standard atmospheric level (approximately 316 parts
     per million).

2.   Net rainfall inputs of H  ion to soil and aquatic systems should
     not be calculated simply as the sum of the H  ions contained in the
     rainwater if pH values above 5.0 are encountered.  Appropriate
     methods would include; H  minus HCO~3, acid anions minus basic
     cations, and net titratable acidity.

3.   Reporting ions input to soil or aquatic systems on a weight per
     unit area basis is often misleading.  Reporting on the basis of
     chemical equivalents per unit area is recommended.

4.   Research should be conducted that will quantify the effect of
     rainfall acidity on soil systems.  This should include the theoretical
     soil chemistry of low level chronic acid inputs on soils, applied
     measurements of acidification rates in laboratory and field systems,
     and methods of evaluating the interaction of acid inputs and plant
     systems on soil acidity.

5.   In conjunction with (4) above systems models should be developed to
     simulate the long-term effects of rainfall acidity on soil and
     soil-piant systems.

6.   The potential acidification resulting from NH 4 inputs should be
     considered in ecological investigations concerned with acid rainfall.

                              SECTION III


     Presently the phenomenon known as "acid rainfall" is arousing a
great deal of interest.  Evidence has been accumulated to support the
assumption that in some areas the acidity of rainfall has been
increasing in recent years as a result of man's activity.  This increase
has been assumed to be largely due to increased atmospheric inputs of
SOp from anthropogenic sources (Likens, 1972; Likens and Bormann, 1974).
Various investigators have called attention to possible adverse effects
of the acidity on terrestrial and aquatic ecosystems.

     Several important aspects of this problem may not have received
adequate attention in the scientific literature.  These include problems
associated with collection, measurement, and reporting, as well as
possible interactions of rainfall acidity with soil-plant processes such
as the cycling of sulfur and nitrogen.  The purpose of this paper is not
to support or to contest the validity of the hypothesized adverse
ecological effects due to acid rainfall.  Rather, the information and
concepts presented here are intended to be useful in the formulation of
testable hypotheses and to provide a firmer basis for future experiments
designed to measure these ecological effects.

                               SECTION IV


     Various investigators have recognized the importance of the carbon
dioxide-bicarbonate system in determining the pH of otherwise neutral
rainwater (Barret and Brodin, 1955; Pearson and Fisher, 1971; Likens and
Bormann, 1974).  Unfortunately many of the reporting practices and
interpretations found in the literature suggest that this system is not
adequately taken into account.  While a dissolution of CaC03 dust by
acidic aerosols could certainly result in bicarbonate in the rainfall as
suggested by Pearson and Fisher (1971), this mechanism is not necessary
to account for the presence of the bicarbonate ion.  Observed
bicarbonate levels of rainfall are best understood by considering the
equilibrium with atmospheric carbon dioxide.  Perhaps a short
explanation will be useful to workers concerned with rainwater chemistry
and its ecological implications.

     The generally accepted standard atmospheric C02 concentration is
approximately 316 parts per million (ppm) or a partial pressure of
3.16 x 10   atmospheres.  The reactions of C0? and water can be
represented by (1) and  (2) below and these are summed to obtain (3).

               C02 + H20? H2C03                                  (1)

               H2C03 i H+ + HC03~                                (2)
               C02 + H20 £ H+ + HC03"                            (3)

     The logarithms to the base 10 of the equilibrium constant  K at 25°C
for (1) and (2) are given by Sillen and Martell  (1964) as -1.46 and
-6.35 respectively.  These are summed to obtain  log  K of -7.81 for (3),
and the equilibrium expression is given by  (4).

               [H+] [HCO "]          y R,
                        3     =   10-7.8i

     The [H ] and [HC03~] are the ion activities in moles per liter and
are considered to be identical to concentration in this very dilute
system.  The [CO^]  refers to the partial pressure of C02 in

     If we restrict the system to atmospheric C09, [CCLl  becomes 3.16 x
  -4                                                    "
10   and (4) can be rearranged as follows:

               [H+] [HC03"3 = (10~7>81)  (3.16 x 10"4)

from which we obtain Equation (5).

               [H+] [HC03"] = 10'11'31                           (5)

     In a pure water system with C0? as the only source of HCO~~ and
                                   £-                          O
assuming sufficient acidity to keep direct hydrolysis of water
insignificant compared to the [H ] ion derived from the C09 equilibrium
               _i_                                          <-
process, the [H ] concentration must equal [HC03~].  Therefore, from (5)
above,  we can write:

     Thus, we expect a pH of 5.65 (and pHCO., of 5.65) for pure rainwater
in equilibrium with atmospheric C0? at 25°C.  The concentration of H
and HCCL~ would be 2.2 x 10~6 moles per liter.  With a rainfall of one
meter annually, the "loading" would be 2.2 x 10   moles per square meter
per year.  Hydrogen ion loading values in this range are not meaningful
as the concentrations within the soil system are controlled by the C02
partial pressure in the soil.  If rainfall or soil solution is more acid
than pH 5.65 for reasons other than the effect of C02, the [H ] term in
Equation (5) increases and the IHCO-,"] decreases.  Bicarbonate
concentrations decrease to insignificant levels between pH 5 and 4, and
in more acid systems C02 partial pressures become unimportant in
controlling pH.

     In basic or near neutral solutions in equilibrium with C0?, the
HC03~ ion is very important.  From the dissociation of water at 25°C as
shown in Equation (6), we obtain Equation (7).

               [H+] IOH] = 10"14                                 (6)


     Substituting (7) into (4) above we have:
                    HO"14] IHCO" ]
                     ~OTT~       J   =  10-7.81
     Which can be rearranged to (8)

               [HC° 3]       ^   106.19                          (8)
               tC02]g [OH]

     Equation (8) is the equilibrium expression for reaction  (9):
               C02 + OH" t HC03"                                  (9)
     On further rearranging Eq. (8) we have Eq. (10):

               [HCO "]
                               , lq
                             106'19L eo]_                       (10)
which at atmospheric C02 concentration of 3.16 x 10~  becomes:
               [HCO ']
               	—   =   490                                 (11)
Equation (11) states that the concentration of HCO," in equilibrium with
atmospheric COp will be 490 times that of OH".  This is a very important
relationship and explains why we would not normally expect to encounter
a "basic rain" phenomenon while acid rain is common.  A likely source of
basic alkaline earth cations in the atmosphere would be the burning of
fuels high in these cations.  These would enter the atmosphere as
oxides— i.e. CaO, Na?0, MgO etc.—and, on encountering water
droplets, would be hydrated to hydroxides.  The hydroxides would be very
basic if it were not for the presence of C0? in the atmosphere.
However, the OH" will be rapidly converted to HCO-," and, at equilibrium,
490 bicarbonate ions will be present for each hydroxyl.  This provides
an effective buffer and prevents highly alkaline conditions from
occurring at the ionic concentrations found in rainfall.





A convenient form of Equation  (5) above is shown  in  02):

               [HCQ-]   =   TQPH-11-31                           (12)

Equation (12) shows clearly that, for a system in equilibrium with
atmospheric CCL at any given pH, the HC03~ concentration is fixed or,
conversely,for any HC03~ concentration the pH is fixed.  If data are
reported that do not fit this relationship the system probably was not
in equilibrium with a C02 partial pressure of 3.16 x 10~  atmospheres.
The graphical presentation in figure 1 over the range usually
encountered in rainfall samples shows clearly that alkaline rainfall
could only occur in the presence of very high bicarbonate levels.

     The above discussion ignores the carbonate ion.  This need not
concern us here, because the mole fraction of dissolved carbon in the
CO,   forms is negligible at the pH values considered.  As pH increases,
C03   increases.  If pH values above about 8.5 were encountered, this
ion would become significant.

     The relevance of the bicarbonate-hydroxyl and bicarbonate-hydrogen
ion equilibria becomes apparent when we consider the overall chemistry
of atmospheric aerosols.  As mentioned above, if bases enter the atmos-
phere from the burning of fossil fuels, particularly coal, they would be
in the oxide form and would hydrate to the corresponding hydroxide.  On
further reacting with acid aerosols formed from SO  and NO  gases, they
                                                  A       A
would be neutralized.  If the acids exceed the bases, the rainfall will
be acid, i.e., the anions will be present in larger amounts than the
basic cations, and electrical neutrality will maintained by the presence
of the H  ions.  The rainfall pH will be below the 5.65 expected from
C02 equilibrium, HCO^  will be very small as shown in Figure 1, and the
well-known acid rainfall phenomenon results.

     If the basic cations exceed the anions, substantial quantities of
   o' will be expected rather than high levels of OH~.  This result can
be predicted from (11), as under equilibrium conditions the HCO," will
be 490 times as high as OH~.  In most cases the HC03~  ion  may  be  regarded
as the alkaline component of rainfall rather than the OH~ ion.  When
rainfall containing HCO-" enters an acid soil, the bicarbonate acts as a
                       •J          +
base, each HC03  reacting with 1 H  ion to form C02 and H^O as shown by
the reverse of equation (3).

     An interesting effect occurs if mean concentrations are calculated
for two or more samplings.  Barret and Brodin (1955) suggest calculating
the H  concentration for each sample and multiplying by the measured
sample volume.  These are summed, divided by the total volume and the
resultant concentration reconverted to a pH value.  This is the usual
method of calculating "average" pH values and H  ion loading.   The
implications are that this would be the pH of the total rainfall
collected if all samples are mixed together.  Unfortunately, if HCO-~
was present in any of the samples this would not be the case,  at least
after equilibration with atmospheric C0?.

     Consider two samples, (a) at pH 7.0 and (5) at pH 5.0.  At
equilibrium with the atmosphere, the HCO~~ concentrations would be 49.0
and 0.49 micromoles per liter for (a) and (b), respectively.    If we
calculate an average H  and HCO^" concentration by the above procedure
for a mixture of equal volumes of (a) and (b), the calculated  pH of the
mixture is 5.30 and the HCO~~ concentration is 24.7 micromoles per
liter.  This solution in actual practice would only be compatible with a
C09 partial pressure of about 6000 ppm.  If we physically mix  (a) and
  <-                                                                    +
(b) and allow it to equilibrate with the atmosphere, equal amounts of H
and HCOQ~ would react to form C09 and HLO until the relationship shown
       -j                        *-                  +
in equation (4) was satisfied.  The fact that the H  and HC03
concentrations would decrease by equal amounts allows us to calculate
the resultant pH and HC03~ as follows.  Let:

                              change in HC03" or H+
               [HC03~]   =    (24.7 x 10"6)-X

By substituting these values into equation (5) and solving the resultant
quadratic, we find that H+ and HC03~ decrease by 4.8 x 10"  moles per
liter during equilibration.  The final pH of the mixture would be 6.6,
and the HC03~ concentration would be 19.9 micromoles per liter.   This
illustrates that if bicarbonates are present, the average pH as  usually
calculated would not be the same as the measured pH of composite
     The implications of this effect are important.  Hydrogen ion load-
ings per unit area are normally calculated by the same procedure as is
average H  concentration or pH.  The above example shows that if two
rainfall events are collected separately the calculated H  loadings may
not be the same as if the two samples are allowed to accumulate in the
collection device prior to analysis.  Also if one event were divided
into several successive sampling periods, and each period was analyzed
separately, the average pH and H  loading observed would be different
than that obtained if the collection accumulated throughout the event.
The possible variations are endless and, in my opinion, indicate a
fundamental error in the usual methods of analysis and interpretation.
     Changes in CCL partial pressure during sample collection and
storage could significantly affect pH and bicarbonate measurements on
many rainwater samples.  For illustration, Figure 2 plots the HC03~ and
pH values for a series of rainwater samples reported by Egner and
Eriksson (1955a, 1955b).  The reported pH values are generally about 0.5


unit lower than would be consistent with a CCL concentration of 316 ppm
but are generally much more consistent with 1000 ppm CO-.  The 1000 ppm
level could easily exist in confined collection vessels or storage
containers, but there is no way to be certain that this is the cause of
the discrepancy.  Granat (1972) has reported a similar and consistent
effect in a large number of samples.  He suggests the effect is due to
the presence of undissolved materials that yield the major ions commonly
found in rainfall upon dissolution, but Granat does not suggest any
compounds with the necessary properties.

     Even though problems of pH and acidity measurements in rainfall
appear formidable, we must address the question of how best to proceed.
In practice few errors will result from pH measurements at unknown C0?
concentrations if the rainfall is consistently acid.   However, if non-
acidic events are interspersed with acid events, the errors could be
substantial .

     The diagram shown in Figure 3 should prove useful  in this regard.
It is developed by utilizing the principle that the net alkalinity b is
independent of C02 partial  pressure, and in systems where the C0~~ ion
is negligible, is defined by (13).

               b = [OH"] + [HCO ~] - [H+]                        (13)

Rearranging (13):

               [HCO ~] = b - [OH"] + [H+]
Substituting for [HCO, ] in (4) above and replacing [OH ] by -^ —
                     3                                       [H ]
we have:
                      (b + [H] - 10")  =  10-7.81

atm  -6
ppm   0
                        log   C02  partial  pressure
         Figure 3
  Diagram showing the effects of varying CC>2 partial pressure on pH.
  Dotted line indicates 316 ppm C02 and numbers refer to pH at 316
  ppm C02-

which rearranges to (14).

               [H+]2 + [H+]b - 10'7'81 [C02] - 10'14 = 0         (14)

For any combination of pH and COp partial pressure, [OH"] and [HC03~]
may be calculated by means of (4) and (6) and b calculated from (13).
To construct figure 3 arbitrary pH values were selected and the values
of b determined for 316 ppm C0?.  The lines were then generated by
              _(.               ^
calculating [H ] as a function of [CO^] using (14).

     No significant effects due to C0? partial pressure variations would
be expected where the lines on Figure 3 are horizontal.  In general the
C0? pressures of interest are in the range of 300 to 1000 ppm, but
higher values may occur.   From the graph we see that with samples below
pH 4.33 no effects occur within the range of interest, but at pH 5.0 C0?
pressures above 1000 ppm would depress measured pH.  With samples above
pH 5.00 the error due to variations in C^ becomes unacceptable.

     Several methods might be employed to minimize possible errors.
First, I suggest that reported pH values should be those taken only
after equilibration with known C0? partial pressures, preferably about
316 ppm.  This would assure uniformity between samples and provide a
uniform basis for comparing data from different investigations.  This
equilibration is particularly important with samples above pH 5.0.
Secondly, H  ion loadings per unit area should not be calculated by
simply determining total  H  from pH and volume measurements.  The long-
term effects on soil systems are more likely to be related to the net H
loading over time.  This would be the difference between H  and
titratable alkalinity or, in practical terms, H  minus HCO-~.  The HCO~~
 represents the alkaline component of the rainfall so in order to

estimate net H  loading from pH measurement, the HC03~ must be
subtracted from the total H+ found.  This method has been used by Granat
(1972).  Alternatively, titration methods could be used and H+ loading
calculated as the difference between titratable acidity and alkalinity.
If all major cations and anions are determined, a method of calculating
net acidification using anion-cation balances as discussed in the next
section should be considered.

                               SECTION V


     An important property of solution systems, whether soil solution or
rainfall, is that electrical neutrality is maintained.  The chemist
analyzing precipitation or stream water utilizes this property when he
checks his results to ascertain whether the cations found equal the
anions.  If the difference exceeds normal  analytical uncertainty, either
an analytical  error has been made or some important constituent has not
been determined.

     This property may be utilized effectively in determining the
acidity of rainfall in the absence of pH measurements, provided the
major cations  and anions have been accurately determined.   Consider a
rainfall  system containing salts of strong acids and  strong bases only.
The pure water system would be approximately neutral and weakly
buffered.  In  the presence of atmospheric C09, the concentration of H
and HCOZ ions  would be approximately 2.6 x 10   moles per liter and the
pH at 25°C would be 5.65 as shown in Section IV above.

     The major anions other than HCOZ involved in the system are NOZ,
  -         -                                             2+    ?+    +
Cl , and SO^ .  Major cations commonly encountered are Ca  , Mg  , Na ,
K , and NH4-  If the anions exceed the cations, electrical  neutrality of
the system is  maintained by H  ions and the system is acid.  This is the
situation we would expect when acidification occurs due to  H9SO. and
     formed from SO  and NO  gases in the atmosphere.    Excluding HCO,~
     ,              X       X                                         O
and H ,  the excess of anions over cations on an equivalent basis may be
used to  estimate net acidity as shown by (15).

               e = 2 [SO/-] + [N03~] + [CT] - 2 [Ca2+]
                   -2 [Mg2+] - [K+] - [NH/]                     (15)

where:  e = excess acid in moles per liter (-e = alkalinity).

     For certain ecosystem effects this measurement may be more
appropriate than direct H  measurements.  For instance, the long-term
effects of soil acidification are probably more appropriately considered
in terms of a deficit of basic cations than in terms of H  ions
directly.  If the excess anions in acid rainfall are mobile, they will
be leached from the soil in association with bases from the exchange
complex, resulting in base depletion and acidification.  Some soils,
particularly those that are acid and highly weathered, exhibit
significant anion exchange properties.  In these soils anion mobility is
reduced, and an excess of acid anions over basic cations may not result
in base depletion.  This condition is discussed more fully in relation
to the sulfur system in the next section.
     If the basic cations in rainfall exceed the anions, the deficit
would be made up by OH" ions in solution if no C0?~ was present.   In the
atmosphere, the OH~ ions are rapidly converted to HCOo" as explained
above.  At any rate, the excess of basic cations over anions (excluding
IICOI) should be an appropriate measure of the ability of the rainwater
to increase the base status of the lithosphere or hydrosphere.
     In practice rainfall acidity measurements derived from anion minus
cation calculation should correspond very closely to values obtained by
substracting HCOo~ from H  derived from pH measurements.  Rainwater at
pH 5.0 contains 10 micromoles H  per liter and 100 micromoles per liter
at pH 4.0.  Cation levels are generally of the same order of magnitude,
so measurements that give acceptable accuracy of the cations and anions
in this range should also combine to give acceptable estimates of net

acidity or alkalinity.  Occasionally rainwaters may be encountered with
ionic concentrations of the order of several thousand microequivalents
per liter.  Cation-anion balance may not give acceptable values for
acidity of the initial rainfall in these cases as the total ion content
is much higher than the H  concentration and relatively small analytical
errors would drastically affect the estimated acidity.  Unfortunately,
when higher concentrations are present, direct pH measurements may also
be a poor indicator of the effect of the water on the acid base
relationships of the soil due to the increased probability of incomplete
ionization of acids or bases.

     In order to examine the validity of this relationship, two sets of
data were selected from the literature and H  ion loading calculated by
the anion minus cation method.  The observed vs.  predicted values were
then compared by linear regression.  For the data of Pearson and Fisher
(1971), the relationship obtained using all 37 points was:
Y = 9.24 + .758 X with a correlation coefficient (r) = .826; where Y is
the observed H minus HCCL in millimoles per square meter per year and X
is the anions minus cations in milliequivalents per square meter.  It
was also noted that a good deal of the scatter resulted from a few
coastal points where high Na  and Cl~ levels appeared to be erratic.
When the coastal points were excluded the least squares regression
equation for the remaining 30 observations was; Y = 5.71 + 0.90 X with a
correlation coefficient (r) of 0.900.  This relationship is shown in
Figure 4.

     A similar comparison was made using data given by Egner and
Eriksson (1955b).  In this case the least squares regression equation
was Y = 0.12  + 0.571 X with a correlation  coefficient (r) of 0.750;
where Y is observed H  minus HCO.,~ and X is calculated from anions minus
cations.  Both X and Y are in units of milliequivalents per square meter
for a one month period and n=43.  One location extremely high in Na  and


C1~ was excluded.  The relationship is shown in Figure 5.  This case is
interesting because most of the values are negative, i.e. the samples
were effectively basic.  The bulk of the scatter appears to arise from
the negative values.  Whether this reflects analytical errors or is due
to other causes is not known.  It is entirely possible that the anion
minus cation values more accurately reflect the net H  loading of the
system than the H  minus HC03~ values.

     Granat (1972) has proposed a model for calculating acidity of
rainfall from the concentration of the individual  ions as follows:

               a = 2 ([SO'l -     - [Na+l + [NO'J - [NH]           (16)
                        g   -

+  -    .    +         2
               b =   ([K+] -    . [Na+]) + ([Mg2+] -      [Na+])       (17)
                     ([Ca2+] -      [Na+])                            (18)
               e = a - 2b

Where:  a = amount of available acid

        b = amount of base expressed as moles of carbonate

        e = excess acid in moles per liter (-e = alkalinity)

This model assumes that all sodium in rainfall is of marine origin.   The
                   2-   +    2+        2+
fraction of the SO.  , K , Mg  , and Ca   assumed to be of marine origin
is equal to the concentration of sodium times the ratio of the
respective ion in sea water to the concentration of sodium in sea water.
The model is apparently intended to represent the net acidity or
alkalinity of rainfall assuming it to be an aqueous solution or
suspension formed from sea salts,  sulfuric acid, nitric acid and

 £    1-0

 I ro
 ^   0-0
                  y  =  -12+-571 x

                  r  =  75O
          ANIONS MINUS CATIONS, meq/m
          Figure 5    Relationship between net acidity as
                     calculated by anions (excluding
                     HC03~) minus basic cations, and
                     observed H+ minus 11003" for data set
                     D502 reported by Egner and Eriksson
                     (1955b).  One coastal station not

ammonia.  Implicit in the model is the assumption that the ratio of Cl~
to Na  is the same in rainfall as in sea water.   Granat shows that the
model is reasonably accurate in predicting the net acidity or alkalinity
of rainfall.
     It may not be readily apparent to the causual  reader, but this
model is actually a form of the strong acid anion minus basic cation
method of calculating acidity shown in equation (15).   The models are
identical if the molar ratio of Cl" to Na  in the rainfall is 1.19, the
ratio found in sea water.  Deviations of acidity or alkalinity
calculated by Granat's model from measured values should be equal to the
difference between measured Cl~ concentrations and an  estimated Cl"
concentration determined by multiplying the Na  concentration by 1.19.

     The concept of cation-anion balance and electrical neutrality of
the system is also useful in considering acid-base relationships of
plant uptake of ions.  In order to maintain electrical neutrality in the
system, the plant must either take up equal amounts of cations and
anions on an equivalent basis or, if the amounts are unequal, it must
release to the solution an amount of ions equal to the difference and of
the same charge as the excess.  A reasonable conceptual model of this
effect is that the plant gives off H  ions equal to the excess of
cations over anions taken up, and indeed this seems to be the case
(Fried and Broeshart, 1967 p. 97).  The converse of this is that an
excess of uptake of anions over cations is balanced by a release of OH~
ions which appear in the media as HCO^ due to the presence of C02-

     Most plants take up more anions than cations, but this may be
reversed if the bulk of the nitrogen uptake occurs as  NhL  rather than
N03".  Variations of a magnitude sufficient to cause major differences
in the acidity relationships of the media may occur.  Healthy plants
usually contain substantially lower quantities of the  usual anions, i.e.

NO^, Cl~, S042", and H2PO.~, than of the basic cations.  This results
from the reduction of S(L   and N0o~ to reduced forms of S and N which
are utilized in the protein structure.  The excess of basic cations are
electrically balanced by organic anions formed in the plant.  During the
breakdown of the dead plant material, these basic salts of organic acids
may be hydrolyzed.  The bases are removed and leached as bicarbonates
and the remaining litter is acidic due to the presence of these organic
acids.  If the pH lowers to the point that bicarbonates can no longer
exist, this mechanism of leaching is no longer operative.  Bases than
can only leach in conjunction with the above anions or by chelation with
certain organic complexes.

     The acid base relationship of this soil-plant system are complex.
It does not seem justified to regard this in terms of a simple input-
output system of  H  ions; the cation-anion balance and acidity
relationships of the whole soil-plant system must be considered.
Apparently one critical research need, in terms of evaluating ecological
effects of acid rainfall, is to synthesize present knowledge of the
acidity and cation-anion relationships into an overall model from which
testable hypotheses can be developed.  Some important aspects of this
system are considered in more detail in the following sections.

                              SECTION VI

     An important aspect of the acid rainfall problem concerns the
possible effects of sulfur inputs from anthropogenic sources on soil-
plant systems.  Most papers seem to focus on the acidity of rainfall,
presumably resulting from sulfuric acid formed by the oxidation of
atmospheric sulfur dioxide.  It would seem desirable to consider
possible effects from the standpoint of the total sulfur cycle, rather
than from a consideration of an isolated portion of the cycle.  Thus a
brief description of the acidity relationships of this cycle appears
relevant and appropriate.

     While sulfur may be present in a variety of forms in plant tissue
and plant residues, it is mostly in reduced (S ") form.   The bulk of the
sulfur is found in the amino acid units of the various proteins.   Micro-
organisms that break down plant residue and soil organic matter may
utilize directly the reduced sulfur in the organic material  or, in some
cases, may utilize sulfate sulfur from the surrounding media.   At any
rate, the organic sulfur compounds are transformed to reduced or incom-
pletely oxidized forms such as sulfides, elemental S, thio-sulfates
etc., (Starkey, 1966).  These reduced or incompletely oxidized  compounds
are then apparently oxidized to sulfate by aerobic chemautotrophs in the
soil.  The oxidation of the reduced sulfur from plant material  and soil
organic matter to the sulfate form will  inevitably be accompanied by the
release of two H  ions for each sulfur atom oxidized.   This  general
reaction may be represented by equations (19)  and (20)  for sulfide and
elemental  sulfur:

               H2S + 2 02 -> 2H+ + S042"                                (19)

               S + 3/2 02 + H20 + 2H+ + S042"                          (20)

     Thus in Figure 6, the oxidation of organic sulfur to SO^   is shown
as releasing 2 H  ions.  H2S may also act as an acid, but the
disassociation constants are such (pK-, = 7.24, pFL = 14.92) that it
would be largely non-ionized in acid soil systems.
     If we assume the SCL ~ released is rapidly taken up by plants, the
acid formed by the oxidation will be balanced by the release of 2 OH
(or HCOo~) ions from the plant in the uptake process (Figure 6).  The
               +                                                      +
net change in H  ions for the overall process in the soil is zero (2 H
and 2 OH~ ions formed), and no change in acidity results.  The consider-
ation of the OH~ ions released by the plant as a result of uptake of
anionc  is an essential point.  However, it cannot be validly considered
in isolation, but must be considered as a part of the total cation-anion
balance of nutrient uptake and plant processes.

     Plants apparently take up sulfur almost entirely as the sulfate ion
but, as noted previously, the preponderance of sulfur in the plant is in
the reduced form.  Sulfate is reduced to the S   form within the plant
prior to its incorporation into the structure of the plant in forms such
as the  sulfur containing amino acids.  This reduction results in the
removal of 2 H   (Figure 6) ions according to a reaction that may be
considered the reverse of (19) above.  The consumption of 2 H  ions at
this point balances the loss of 2 OH" ions, (which may be considered as
equivalent to formation of 2 H  ions) during the uptake of sulfur.

     The total H  ion production and consumption of the sulfur cycle
(Figure 6) is balanced and no net change in acidity should result.
While the cycle represents a drastic simplification of the sulfur
oxidation-reduction processes in the system, the acidity relationships
should be basically valid.


                                    on aerobic
                            Figure 6    Simplified sulfur cycle showing acid
                                      formed or consumed by the various

     The anaerobic reduction of sulfate to sulfide (Figure 6) requires
low oxygen levels and a fixed carbon energy source, as it is carried out
by heterotrophic anaerobic organisms.  High concentrations of sulfide
compounds may accumulate in flooded soils where sulfate is continually
replenished, such as river deltas.  This process tends to produce
neutral or basic soils as H  ions are consumed.  Drainage of these soils
results in rapid acidification due to H  produced during the oxidation
of S2" to S042".

     While the sulfur cycle is inherently balanced as far as acid
production and consumption are concerned, it does provide some potential
for leaching of bases and subsequent development of soil acidity.  This
arises because lags may occur between oxidation to sulfate and uptake of
the sulfate ion.  If sulfate ions accumulate, there is a concomittant
increase in H  formed in the oxidation of reduced sulfur to sulfate.
The H  ions replace basic cations on the clay and organic matter
complexes, and these bases may be leached from the soil in association
with the SCL   anion.  The quantification of any acceleration of this
loss of cations that may be due to acid rainfall is a critical aspect of
the assessment of possible ecological effects that may result.
Complications arise in this quantification due to the
inter-relationships between the nature of soil acidity and the process
of sulfate adsorption in soils.
     In neutral or acid soils sulfate ions may remain in the soil
solution or become adsorbed on soil particles.  Various conceptual
models have been proposed but for our purposes it is most useful to
consider this adsorption to be an anion exchange phenomenon.  Soil
particles are generally negatively charged but some positive charges
exist as the result of the amphoteric nature of soil organic matter and
broken bonds on the edges of the clay lattices.  As soils become acid,
the alumina silicates tend to decompose as represented by equations 21-

               A1(OH)3 + H+ %• A1(OH)2+ + H20                         (21)
               A1(OH)2+ + H+ %- AL3+ + H20                            (23)

     According to this model most reserve acidity in soils does not
consist of H  ions adsorbed on negatively charged soil colloids but
exists in the form of monomeric Al   and positively charged hydroxy
aluminum complexes and polymers.  These positive charges neutralize
negatively charged colloids or adsorbed anions such as SO^  .   As soils
become more acid their cation exchange capacity decreases due to
neutralization of negative charges by this mechanism.  The increase in
positive charges with acidification increases the capacity of soils to
adsorb anions such as sulfate (Harward, Chao, and Fang, 1962;  Harward
and Reisenaur, 1966).  Apparently as acidity increases due to anthro-
pogenic sulfur inputs, the ability of the soil to hold the sulfate ion
through the adsorption mechanism would also increase.

     Recent data of Nyborg et. al . , (1973) indicate that direct soil
absorption of S02 may be substantial.   Directly downwind from a source
in central Alberta, they found enrichment in the range of 10-20
kilograms S per hectare over a three month period.

     If reduced sulfur such as elemental S° is added to an aerobic soil
system, we can expect oxidation to SO,   by soil organisms and the
                        +            ^
concomitant release of H  ions.   If acid rainfall containing H?SO»
enters the system, the H  ions are supplied directly.  If the rainfall
contains a mixture of H2SO» and H2$03 (Brosset, 1973), the pH of the
rainfall would not be as low as for a similar concentration of HUSO.
alone, due to the lower dissociation of H?SOo-  However, as soon as

temperature and moisture conditions become favorable, the SCL   would
undoubtedly be oxidized to SO-   in the soil.  Thus, the net effect on
the acidity of the soil system would be the same of H2$03 as for H2S04.

     From the above we see that the net increase in acidity per mole of
sulfur added is identical for reduced sulfur, sulfuric acid or sulfurous
acid.  If acid rainfall containing H^SCL enters the system, H  ions are
supplied directly.  The dissociation constants for H^SO., are much lower,
but added sulfite would be oxidized to sulfate assuming temperature and
moisture levels are favorable.  In all cases, the acidity is increased.
However, if the SO,   is rapidly absorbed by the plant systems, the
resultant release of OH" ions (Figure 6) would tend to balance the extra

     Sulfur loads calculated from the data of Pearson and Fisher (1971)
for the northeastern U.S., tend to be in the range of 8 to 15 kg ha~
yr~ .  Jonsson and Sundberg (1970) give a value of 9 kg ha"  yr~  for
the Swedish forests.  Values of this magnitude could probably easily be
absorbed by most agricultural ecosystems as they are in the same general
range as harvest removals of sulfur.  In forest ecosystems a long-term
build up in soils, plants, and litter might be expected.  If loading is
greater than the capacity of the ecosystem to utilize or store sulfur,
leaching losses will probably occur, bases will be lost in conjunction
with the SO,  , and the soil will  become more acid.

     The above discussion refers specifically to inputs of sulfur that
are not balanced by basic cations.  If rainfall inputs of sulfate are
                                                         +    2+    2+
accompanied by similar amounts of basic cations, i.e., Na , Ca  , Mg  ,
or K ,  little change in soil acidity would be expected.   Over long
periods of time, however, if the distribution of these input ions is

different than the distribution of ions in the soil  solution and in
equilibrium with exchangeable cations in the soil,  a shift in the
cation equilibria could occur.   The effect of such  a shift would depend
on the resultant distribution.

     The capacity to utilize additional sulfur is often closely related
to the nitrogen cycle.   Nitrogen sulfur ratios of most plant proteins
will be in the range of 12:1 to 16:1, while ratios  in soil  organic
matter tend to be of the order  of 8:1 or 10:1  (Stewart and Whitfield,
1965).  When nitrogen being taken up by plants is largely derived from
soil organic matter, the growth limiting nutrient is more likely to be
nitrogen than sulfur.  The sulfur mineralized will  probably be sufficient
for utilization of the nitrogen mineralized.   However, if nitrogen comes
from an outside source such as  fixation of atmospheric nitrogen or non-
sulfur containing fertilizers,  the probability of sulfur deficiency is
substantially increased.  Thus  sulfur responses in  agriculture are
generally found on legume crops such as alfalfa and clover or on crops
where heavy nitrogen fertilizer applications are made using fertilizer
materials that do not contain sulfur.  In forest ecosystems where
nitrogen fixation is rapid such as in alder stands,  there is a definite
possibility that anthropogenic  sulfur inputs would  result in accelerated
growth.  In such systems the removal of sulfur as a growth limiting
factor would probably increase  productivity.   It is even possible that
sulfuric acid in rainfall would not result in increased soil acidity due
to the balancing effect of the  OH~ ions given off by the plant in the
uptake of SO,  .

     Mature timber stands or other systems in which nitrogen is supplied
largely from internal cycling are less likely to be sulfur limited.  If
the capacity of these ecosystems to utilize sulfur  inputs is exceeded,
the probability of deleterious  effects from sulfuric acid containing
rainfall  would be increased.

     Finally, the effect of sulfur inputs on base rich soils should be
considered.  Such soils are generally found in arid areas, but occasion-
ally occur under humid or subhumid conditions.  These soils are
virtually 100% base saturated, i.e., all cation exchange positions are
occupied by bases, and they generally contain free alkaline earth
carbonates.  The acid-base relationships of the sulfur cycle here are
identical to those discussed above except for the effect of excess acid.
Free carbonates result in a large buffering capacity.  Acid rainfall
high in sulfate ion would dissolve the carbonates, and sulfates would
tend to accumulate at some point in the profile and precipitate out as
gypsum.  This should not cause a problem as many soils naturally contain
layers where free gypsum crystals are present.  Sulfur bearing fuels
might well be exhausted before large areas of soil carbonates are
depleted.  However, some ecological effects are possible as many
calcareous grassland soils are actually acidic near the surface.  Acid
inputs from acid rainwash might be expected to intensify and extend this
layer.  While the process would probably be very slow, the long-term
ecological consequences are unknown.

                              SECTION VII


     In light of the recent interest in the effect of anthropogenic
sulfur inputs on the acid-base relationship in soil-plant systems, it is
appropriate to call attention to the role of the nitrogen system in this
regard.  From an overall ecosystem point of view, the nitrogen cycle can
be expected to release or consume much greater quantities of H  ions
than would be involved in the sulfur cycle.  Also the contribution of
nitrogen compounds to the potential of rainwater to acidify soil systems
may be much greater than generally recognized.  Plant uptake of nitrogen
occurs in both the NH» and NOo" forms.   Factors that determine the
relative uptake in each of these forms include the nature of the plant
and the presence  or absence of conditions favorable to the oxidation of
NH,  to NO.," in the soil system.

     While plants may contain many types of nitrogen compounds, the
major portion of the nitrogen is found in the structural proteins.   This
nitrogen is in the N   oxidation state,  the same as is found in NH->,
   +                                                              J
NH»  and soil organic nitrogen.   When plants take up nitrogen in the
NH,  form, no change in oxidation state  in the system is required
(Figure 7).  Starting with the soil organic component in Figure 7,  we
consider the release of N to be a simple deamination of -CNH? to NH~
with no change in acid-base status required.  This NH~ will, require 1
 +                     +
H  ion to ionize as NH. , or if the process is considered to be a
hydrolysis, one OH" is formed (Equation  18).

               NH3 + H20 -v*" NH4+ + OH"                                (24)

     The uptake of NH^  by a plant results in the release or exchange of
an H  ion so the net change in acidity for the soil system is zero.  If
ammonification occurs without either uptake or oxidation to NOo so that


R -•« M H -^ . M
*\ J
^ organic nitrogen H
R — C. — R 	 	 ^- M H — — -^ 	 »— M

                                    R  organic  nitrogen
                                     Figure 7     Simplified nitrogen cycle assuming  all
                                                 N  utilized in the NH^ form, and showing
                                                 acid formed or consumed by the various

NH.  accumulates, we might expect the pH of the system to increase.
This may occur but the relationship is complicated by the basic nature
of the original -NhL group and by the formation of keto groups in the
usual oxidative deamination.  The acid-base relationships in the plant
are relatively simple when nitrogen is taken up in the NH»  form.  One
 +                                                                    +
H  is released in the conversion to NH3 or C-NH,,.  This balances the H
ion released in the uptake process so the net change is again zero.  The
acid-base relationship of the cycle appears to balance as one would
intuitively expect.

     The system becomes more complex when we consider the oxidation of
NhL  to NO," and the subsequent uptake of the nitrate ion (Figure 8).
Starting again with the breakdown of soil organic matter, we represent
the system as releasing NFL.  The subsequent protonation of NH, to form
   +             +                           +       -
NhL  requires 1 H  ion.  The oxidation of NhL  to N0?  and subsequently
   _                        +
NO, , releases a total of 2H  ions, as shown below:

               NH4+ + 3/2 02 + 2H+ + N02" + hU.0                       (25)

N02" + 1/2 0
NH4+ + 2 02
2 + N03~
+ 2H+ + N03" + H20
     These transformations are carried out in the soil largely by
chemautotrophs, step (a) by organisms of the genus Nitrosammonas and
(b) by Nitrobacter.  The subsequent uptake of NO," releases 1 OH~,
which balances one half the acidity released in the oxidation.  The
other half is balanced by the H  utilized in the protonation of NH,.
Again, we find the net change in the soil is zero and the system
balances, provided the cycle is completed by plant uptake.

            organic  nitrogen
          i 2

         R organic nitrogen
          Figure 8    Simplified nitrogen cycle assuming all II utilized in the
                     form, and showing acid formed or consumed by the various

     The plant transformations may be regarded as the reverse of those
found in the soil.  Two H  ions are consumed in the reduction of N(L~ to
NHL   by the nitrate reductase enzyme system.  One of these is balanced
  ^                +
by the loss of an H  to form NH3 prior to the amination.   The other is
balanced by the release of the OH" that occurred during the uptake of
NO,", as an OH" lost by the plant is equivalent to the formation of an
     Even though the overall system can be considered balanced, the
nitrogen cycle may, under certain conditions, offer potential mechanisms
for losses of bases and subsequent  acidification of the soil system.
When mineralization occurs followed by the oxidation of NH.  to nitrate,
     4.                                                    ^
the H  ions formed will replace a basic cation on the cation exchange
complex.  This basic cation is then subject to leaching in conjunction
with the N03~ ion.  If conditions allow any build-up of NO.,", the poten-
tial for leaching exists when water passes through the profile.  Most
natural ecosystems tend to maintain low nitrate levels as the nitrates
formed are rapidly taken up by the plants.

     Some interesting questions arise when one considers the potential
acidity relationships of mineral nitrogen inputs in precipitation.  Any
anion taken up by a plant will cause the release of an OH~ ion, so
nitrate could be considered as resulting in a basic reaction.  The
highly mobile nitrate ion is rapidly absorbed by the plant, and Pearson
and Fisher (1971) note that stream loads of mineral N are less than
precipitation inputs.  Even though this possible basic effect may be
valid, one must be cautious in the interpretation of this effect in
isolation from the rest of the cation-anion system.
     Perhaps the most interesting aspect of the nitrogen cycle in
relation to rainfall acidity involves the ammonium ion.  If the NH,
falling in precipitation is oxidized in the soil to the nitrate form,

2 moles of H+ will be formed for each mole of NH4  oxidized as shown
by equation (27).  The data of Pearson and Fisher (1971) show that this
effect may be substantial.  Table 1 shows values of the mean ionic
concentration and loading for Mays Point, New York in 1966, calculated
from their data.  The NH4+ value of 15.1 milliequivalents of NH4  per
square meter at this location also happens to be the mean value for all
locations and years that they report.  At this location they report an
H  loading of 34.9 milliequivalents per year which would correspond to a
mean value of 44.8 microequivalents per liter with a total  precipitation
of 78 centimeters.  This mean H  concentration results in an "average
pH" of 4.35.  If 2 moles of H  are released in the oxidation of each
mole of NH4+ to N03" (Equation 27), the oxidation of 15.1
milliequivalents of NH.  to NCL" will result in the release of 30.2
milliequivalents of H  per square meter, an amount very similar to the
34.9 milliequivalents reported as the direct input from H  loading.
     Another example of the possible significance of the effect of
acidification due to oxidation of NH.  can be seen from data shown by
Likens (1972) and Likens and Bormann (1974).  These authors call  atten-
tion to an apparent shift from NHL  to N03~ as the dominant form of
inorganic nitrogen at Geneva and Ithaca, New York since about 1945.
They also note increasing acidity of rainfall over the same period as
evidenced by pH values of about 4.0 at present.   No direct pH measure-
ments are available for the early samples, but they are inferred to be
higher due to the recorded presence of bicarbonates and from their
reaction to methyl orange indicator.  The data are presented in
graphical form, but apparently a value of 0.9 milligrams NH^  per liter
is representative of the bulk of the observations prior to 1945,  with a
few values above 3.0.   The release of 2 H  ions  per mole on oxidation of
0.9 milligrams NH,  per liter would release 0.106 milligrams H  per
                 ^                       +
liter.  This is almost identical to the H  concentrations, shown by


                              Milligram                Microgram
                             Equivalents,             Equivalents,
     Cations                meter"2, year"1             liter"1

     Ca2+                          35.7                     45.9
     Mg2+                          11.7                     15.0
     Na+                            8.1                     10.4
     K+                             3.1                      4.0
     NH4+                          15.1                     19.4
     H+                            34.9                     44.8

     S042"                         87.9                    113.0
     Cl~                           10.5                     13.5
     NO ~                           2.1                      3.5
*Data as shown by Pearson and Fisher was in units of tons per square
mile per day.  Precipitation totaled 78 centimeters.

Likens (1972) for six collection stations near Ithaca, New York in 1970-
71.  This, of course, may not infer that if oxidation of the NH^  in
these early samples had occurred, the H  concentration would have been
0.106 milligrams per liter (pH = 3.98) as the buffering effects of
bicarbonates in the system cannot be estimated.  However, a
consideration of the acidification potential of the ammonium present may
well modify substantially our conclusions concerning the relative
acidification effects of the early rainwater samples as compared to
present day "acid rainfall."

     It is also important to recognize that the oxidation of NH,  to
NO-  by aerobic chemautotrophs may occur in rainfall collection vessels
unless appropriate precautions are taken.  If this occurs acidification
results.  We can see by the above examples that the potential increase
in observed acidity could be significant; in the case shown in Table 1,
for instance, the oxidation of the NH»  to NO.," would drop the "average
pH" from 4.35 to 4.03, nearly doubling the H  loading as usually calcu-
lated from pH measurements.  Serious questions about the validity of
many pH measurements of rainfall samples might be appropriate on this

     One could argue that the appropriate pH would be that taken after
oxidation occurs, as the acidification effect will be manifested when
the oxidation occurs in the soil.  This has the disadvantage of ignoring
the acid-base relationships of the nitrogen uptake processes.  We seem
to be forced to conclude that the acidification effects of rainfall are
closely linked to other ecosystem processes, particularly those of the
nitrogen cycle.  Consideration of the acidity of rainfall in isolation
may be an unwarranted oversimplification.
     When ammonium is taken up directly by the plant, acidification
still occurs due to the H  given off by the plant in the uptake process.
The magnitude of this effect is often underestimated.  In the past many

experiments showing poor growth in solution and pot cultures when ammon-
ium was used as a nitrogen source were interpreted as a toxic effect of
NH»+ ion when in fact the toxicity was caused by acidity developed in
  "                                                     +
the medium.  Carnation producers had long considered NH^  as an unsuit-
able source of nitrogen, but it has been shown that they do well when
NH/ is used if CaC03 is added to prevent acidification (Schekel, 1971).

     Ammonium inputs, whether oxidized to nitrate or not, can thus be
shown to have an acidifying effect.  While it is probably extreme to
consider the total potential acidification as representing two moles H
per mole of NH/1", the effect should be recognized in our evaluation of
the effect of rainfall acidity.

     If ammonium inputs are important, it may be well to consider their
source and distribution.  In past years, many measurements of nitrogen
inputs in rainfall were made for the purpose of evaluating their agri-
cultural significance.  Results vary widely, and it is impossible to
accurately evaluate whether the variation is due to contamination and
analytical inaccuracies or actual variation in amounts.  It has been
noted that the results quite consistently show a 2:1 ratio of ammonium
to nitrate N (Stevenson, 1965), but the results of Pearsons and Fisher
(1971) show widely varying ratios both between locations and between
years at a single location.

     Atmospheric ammonia levels vary widely.  Sources of ammonia often
widely quoted include industrial, atmospheric fixation (electrical and
photochemical), and volatilization from land surfaces.  Recently Luebs,
Davis, and Laag (1973) have shown much higher atmospheric concentrations
of NH~ in a dairy area in California than in a control area.  They also
     0     +
claimed NH^  inputs in the precipitation in the dairy area to be three
times that of the control area.  Hutchinson and Viets (1969) have

reported high ammonia absorption by traps near feedlot operations.
Thus, the possible effects of high concentrations of livestock must be
considered as well.

     Atmospheric NH., reacting with sulfuric acid aerosols would form
ammonium sulfate and bisulfate salts as shown in Equations (28) and

               H2S04 + NH3 + NH4+HS04                                 (28)
                 4+HS04 + NH3 + (NH4+)2 S04                           (29)
     The resultant solution would be much less acidic than aerosols not
neutralized by NHL, yet the system would retain the capacity to acidify
soils as shown by the mechanisms discussed above.  Therefore, the
neutralization of acid aerosols by atmospheric NH- may not reduce the
capacity of the rainfall to acidify soils.

     In conclusion, it probably is not possible at the present time to
predict quantitatively the effect of nitrate and ammonium on the
capacity of acid rainfall to acidify the soil system.  On theoretical
grounds we could expect substantial effects, and evaluation of this
contribution should be included among the goals of future research on
the acid rainfall problem.

                             SECTION VIII


Barret, E. and G. Brodin.  1955.   The acidity of Scandinavian precipita-
     tion.  Tellus VII(2):251-257.

Brosset, Cyril 1.  1973.  Air-borne acid.  Ambio.  11(1-2):2-9.

Egner, H. and E. Eriksson.  1955.  Current data on the chemical composi-
     tion of air and precipitation.  Tell us VII(1):134-139.

Egner, H. and E. Eriksson.  1955b.   Current data on  chemical composition
     of air and precipitation.   Tellus VII(2):267-271.

Fried, M. and H. Broeshart.  1967.   The soil-plant system in relation to
     inorganic plant nutrition.   Academic Press:  New York and London.
     358 p.

Granat, L.  1972.  On the relation  between pH and the chemical
     composition of precipitation.   Tellus XXIV(6):550-560.

Harward, M. E., T.  T. Chao, and  S.  C. Fang.  1962.  Soil properties and
     constituents in relation to  mechanisms of sulphate adsorption.  In:
     Radioisotopes in Soil-Plant  Nutrition Studies.   International
     Atomic Energy Agency, Vienna.

Harward, M. E. and H. M.  Reisenaur.  1966.  Reactions and movement of
     inorganic sulfur.   Soil  Science 101:326-335.

Hutchinson, G. L. and F. G. Viets, Jr.  1969.  Nitrogen enrichment of
     surface water by absorption of ammonia volatilized from feedlots.
     Science 166:514-515.

Jonsson, Bengt and Rolf Sundberg.  1970.  Has the acidification by
     atmospheric pollution caused a growth reduction in Swedish forests?
     Institutionen for Skogsprodukrion (Department of forest yield
     research), Rapporter och Uppsates (Research Notes), Skogshogskolan
     (Royal College of Forestry), Stockholm.  48 p.

Likens, G. E.  1972.   The chemistry of precipitation in the central
     finger lakes region.  Tech. Report No. 50.   Cornell University
     Water Resources  and Marine Sciences Center.  Ithaca, New York.   47

Likens, G. E. and F.  H. Borman.  1974.  Acid rain:  A serious environ-
     mental problem.   Science 184:1175-1179.

Luebs, R. E., K. R. Davis, and A. E.  Laag.  1973.  Enrichment of the
     atmosphere with  nitrogen compounds volatilized from a large dairy
     area.  Jour. Environ. Quality 2(1):137-141.

Nyborg, M., and McKinnon, Allen Associates.  1973.  Atmospheric sulfur
     dioxide:  Effects on the pH and  sulfur content of rain and snow;
     addition of sulfur to surface water, soil,  and crops; acidification
     of soils.  In:  Proceedings of a Workshop on Sulphur Gas Research
     in Alberta (D. Hocking and D. Reiter, eds.).  Information Report
     NOR-X-72, Northern Forest Research Centre,  Edmonton, Alberta,   pp.

Pearson, F. G., Jr. and D.  W.  Fisher.   1971.   Chemical  composition of
     atmospheric precipitation in the northeastern United States.  U.S.
     Geological Survey, Water Supply paper 1535-p.  U.S. Government
     Printing Office, Washington, D.C.  23 p.

Schekel, K. A.  1971.  The  influence of increased ionic concentrations
     on carnation growth.   Jour.  Amer. Soc.  Hort. Sci.  96(5):649-652.

Sillen, L.  G. and A.  E. Martell.   1964.  Stability constants of metal-
     ion complexes.  Ind.  Ed.  Special  Publications, The Chemical Society,
     London.  754 p.

Starkey, Robert L.  1966.   Oxidation and reduction of sulfur compounds
     in soils.  Soil  Science 101(4):297-306.

Stevenson,  F. J.  1965.  Origin and distribution of nitrogen in soil.
     In:  Soil Nitrogen (W.  V. Bartholomew and F. E.  Clark,  eds.).
     Amer.  Soc. Agron.  Mon.  No. 10.  Madison, Wisconsin,  p. 1-42.

Stewart, B. A. and C. J.  Whitfield.  1965.   Effects of  crop  residue,
     soil  temperature,  and  sulfur on the growth of winter wheat.  Soil
     Sci.  Soc. Amer.  Proc.   29:752-755.

                                   TECHNICAL REPORT DATA
                            (Please read Instructions on the reverse before completing)
 \. ntPORT NO.

 Chemical/Biological Relationships Relevant  to
 Ecological Effects  of Acid Rainfall

                                                           3. RECIPIENT'S ACCESSIOI*NO.
             5. REPORT DATE  February 1975
               Date  of  Preparation
 John 0, Reuss
 National Ecological  Research Laboratory
 National Environmental Research Center
 200 SW 35th St.
 Corvallis, Oregon  97330
             10. PROGRAM ELEMENT NO.
             11. CONTRACT/GRANT NO.
                                                           13. TYPE OF REPORT AND PERIOD COVERED
                                                           14. SPONSORING AGENCY CODE
 16. ABSTRACT This paper deals  with problems of measurement  and interpretation of rainfall
 acidity in terms of  effects on the soil-plant system.
      The theory of the  carbon dioxide-bicarbonate  equilibria and its effect on rainfal
 acidity is given.  The  relationship of a cation-anion balance model of acidity in
 rainfall to plant nutrient  uptake processes is discussed,  along with its relationship
 to a model previously proposed in the literature.  Average H  concentration calculated
 from pH measurements  does not appear to be a satisfactory  method of determining H
 leading from rainfall if  the rain is not consistently acid.   Calculating loading from
 H  minus HCO  , strong  acid anions minus basic cations,  or net titratable acidity is
      The flux of H   ions  due to plant uptake processes and sulfur and nitrogen cycling
 is considered.  H  is produced by oxidation of reduced sulfur and nitrogen compounds
 mineralized during decomposition of organic matter.  Plant uptake processes may re-
 sult in production of either H  or OH  ions.  Fluxes of  H   from these processes are
 much greater than rainfall  H  inputs, complicating measurement and interpretation of
 rainfall effects.  The  soil acidifying potential due to  the  oxidation of the NH,
 in rainfall is apparently of a similar magnitude to the  direct acidity inputs.
      This report was  submitted by the National Ecological  Research Laboratory under
 the sponsorship of the  Environmental Protection Agency.  Work was completed as of
 January 1975.	
                                KEY WORDS AND DOCUMENT ANALYSIS
 Water analysis
 Water chemistry
 Soil Science
 Soil Chemistry
 Plant Nutrition
Rainfall  Chemistry
Precipitation (meterology)
Acid  rainfall
Soil  acidification
                                              19. SECURITY CLASS (This Report)
                           21. NO. OF PAGES
                                              20. SECURITY CLASS (This page)
                          22. PRICE
EPA Form 2220-1 (9-73)