PB86-179678
Dissolved Oxygen and
Oxidation-Reduction
Potentials in Ground Water
Illinois State Water Survey Div., Champaign
Prepared for
Robert S. Kerr Environmental Research Lab.
Ada, OK
Apr 86
I
1
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EPA/600/2-86/042
April 1986
DISSOLVED OXYGEN AND OXIDATION-REDUCTION POTENTIALS IN GROUND WATER
by
Thomas R. Holm
Gregory K. George
Michael J. Barcelona
Illinois State Water Survey
Champaign, Illinois 61820-7407
CR-811477
Project Officer
Bert E. Bledsoe
Processes and Systems Research Division
Robert S. Kerr Environmental Research Laboratory
Mi, Oklahoma 74820
ROBERT S. KERR ENVIRONMENTAL RESEARCH LABORATORY
OFFICE OF RESEARCH AND DEVELOPMENT
U.S. ENVIRONMENTAL PROTECTION AGENC?
ADA. OKLAHOMA 74820
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TECHNICAL REPORT DATA
(Plftse read /nitrucliont on the rtvene be fort completing)
1. REPORT NO.
EPA/600/2-86/042
2.
3. RECIPIENT'S ACCESSION NO.
PM6179678/AS
4. TITLE ANDSUBTITLE
DISSOLVED OXYGEN AND OXIDATION-REDUCTION POTENTIALS
IN GROUND WATER
5. REPORT DATE
April 1986
«. PERFORMING ORGANIZATION CODE
AUTHORISE
Thomas R. Holm, Gregory K. George and
Michael J. Barcelona
•. PERFORMING ORGANIZATION REPORT NO
PERFORMING ORGANIZATION NAME AND ADDRESS
Illinois State Water Survey
2204 Griffith Drive
Champaign, Illinois 61820-7407
10. PROGRAM ELEMENT NO.
CBPC1A
11. CONTRACT/GRANT NO.
CR-811477
2. SPONSORING AGENCr NAME AND ADDRESS
Robert S. Kerr Environmental Research Laboratory
Office of Research and Development
U.S. Environmental Protection Agency
Ada, Oklahoma 74820
13. TYPE OF REPORT AND PERIOD COVERED
1nal (June 1984 - March 1986]
14. SPONSORING AGENCY CODE
EPA/600/15
5. SUPPLEMENTARY NOTES
6. ABSTRACT
Water samples were collected from various depths 1n a pristine sand and gravel
water table aquifer at monthly intervals over a period of one year. Dissolved oxygen
concentrations were near saturation 9 feet below tha water table and decreased to
nearly zero at 78 feet below the water table. Changes 1n the Eh values were consisten
with changes in the dissolved oxygen concentrations. Hydrogen peroxide was detected
in nanomolar concentrations at all depths, but not on every sampling run. Of all
oxidation-reduction potentials calculated from substituting analytical results Into
the Nernst equation, only the Fe^/Fe2"1" couple in the deepest well agreed with the
measured Eh within 50 millivolts. For the 02/Ha02 and NOa'/WV couples the range of
potentials calculated from one year's data overlapped the range of measured Eh values,
so there was some agreement on the average. However, for a given sampling run, the
various calculated potentials spanned several hundred millivolts, which means that
the activity ratios for the various couples differed by many orders of magnitude (I.e.
the system was not 1n redox equilibrium). The concentration profiles of many, solutes,
including dissolved oxygen, suggest mixing of shallow and deep ground waters. The
observed concentration profiles were relatively constant over the duration of the
sampling.
KEY WORDS AND DOCUMENT ANALYSIS
DESCRIPTORS
b.tDENTIFIERS/OPEN ENDED TERMS C. COSATI Field/Group
19. SECURITY CLASS (Ttiti Report!
UNCLASSIFIED
21. NO. OF PAGES
DISTRIBUTION STATEMENT
RELEASE TO THE PUBLIC
20. SECURITY CLASS (Thitpaffl
UNCLASSIFIED
22. PRICE
EPA F*m 2220-1 (*•». 4-77) PREVIOUS COITION i* OMOLKTK
1
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DISCLAIMER
The information in this document has been funded wholly or in part by
the United States Environmental Protection Agency under Cooperative Agreement
No. CR-811A77 to Illinois State Water Survey. Although it has been subjected
to the Agency's peer and administrative review and approved for publication
as an EPA document, mention of trade names or commercial products does not
constitute endorsement or recommendation for use.
ii
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FOREWORD
EPA Is charged by Congress to protect the Nation's land, «lr and water
systems. Under a mandate of national environmental laws focused on air and
water quality, solid waste management and the control of toxic substances,
pesticides, noise and radiation, the Agency strives to formulate and imple-
ment actions which lead to a compatible balance between human activities and
the abilir.y of natural systems to support and nurture life.
The Robert S. Kerr Environmental Research Laboratory Is the Agency's
center of expertise for investigation of the soil and subsurface environ-
ment. Personnel at the Laboratory are responsible for management of research
programs to: (a) determine the fate, transport and transformation rates of
pollutants in the soil, the unsaturated zone and the saturated zone of the
subsurface environment; (b) define the processes to be used in character-
izing the soil and subsurface environment as a receptor of pollutants;
(c) develop techniques for predicting the effect of pollutants on ground
water, soil and Indigenous organisms; and (d) define and demonstrate the
applicability and limitations of using natural processes, indigenous to the
soil and subsurface environment, for the* protection of this resource.
Currently, there is a very limited amount of information and under-
standing of oxygen and oxidation-reduction processes in the subsurface
environment of ground water. This report is a first attempt at defining
and understanding the dynamic relationships involved in this environment.
These relationships should be useful to olcroblologlsts, geochemists and
engineers studying ground-water quality and the fate of contaminants In
ground water*.
Clinton W. Hall
DltWOT
Robert S. Kerr Environmental
Research Laboratory
iii
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ABSTRACT
Water samples were collected from various depths in a pristine sand and
gravel water table aquifer at monthly intervals over a period of one year.
Dissolved oxygen concentrations were near saturation 9 feet below the water
table and decreased to nearly zero at 78 feet below the water table. Values
of Eh reflected the dissolved oxygen concentrations. Hydrogen peroxide was
detected in nanoroolar concentrations at all depths, but not on every
sampling run. Of all oxidation-reduction potentials calculated by substi-
tuting analytical results into the Nernst equation, only the Fe3*/Fe2*
couple in the deepest well agreed with the measured Eh within 50 millivolts.
For the 02/H2°2 and N03~/NHjj* couples the range of potentials calculated
from one year's data overlapped the range of measured Eh values, so there
was some agreement on the average. However, for a given sampling run, the
various calculated potentials differed by several hundred millivolts, which
means that the activity ratios for the various couples differed by many
orders of magnitude (i.e. the system was not in redox equilibrium). The
concentration profiles of many solutes, including dissolved oxygen, suggest
mixing of shallow and deep ground waters. The observed concentration
profiles were relatively constant over the duration of the sampling.
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CONTENTS
Page
Abstract iv
Acknowledgements vi
Introduction 1
Conclusions ......... 3
Recommendations 3
M«i*rl«lM and Molhodn .......... D
Well Drilling and Installation H
Water Sampling 5
Analytical Methods 6
Hydrogen Peroxide 7
Results and Discussion 8
Dissolved Oxygen, Trace Metals, and Eh 9
Hydrogen Peroxide 10
Other Solutes 11
Calculating Redox Potentials 13
References 1l|
Appendices 21
A. Oxygen Diffusion Through Sampling Tubing 22
B. Tabulation of Ground Water Analyses 23
Figures 35
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ACKNOWLEDGEMENTS
The authors thank the reviewers for their comments on the manuscript,
the Illinois Department of Conservation for access to the field site, the
Illinois Department of Transportation for well drilling, Ms. Para Beavers for
typing the manuscript, and Ma. Lynn Weiss for drawing figures 1, 2, and 22.
vl
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IHTRODUCTION
Because ground water makes up a significant fraction of the water
resources of the United States, comprising approximately 40J of the water
withdrawn for agricultural, municipal, and industrial use [Anon. 1978],
ground-water quality is of great concern. If ground-water quality is
defined in terms of the concentrations and properties of substances
dissolved in the water, then the presence or absence of dissolved oxygen
affects many aspects of ground-water quality. The mobilities, reactivities,
and toxlclties of many elements may all depend on their oxidation state.
Microblal populations are distinctly different in oxlc and anoxlc waters
and, therefore, the rates of nicrobial degradation of organic compounds are
also quite different. Thus, dissolved oxygen affects both the geochemlcal
and microblal processes which are likely to influence water quality. In
this project we have begun the chemical characterization of a shallow
aquifer with dissolved oxygen concentrations ranging from near-saturation to
neai—anoxic, conditions that are typical of many aquifers which are
susceptible to contamination. The results of this study should, therefore,
be useful to mlcrobiologists, geochemlsts, and engineers studying
ground-water quality and the fate of contaminants in ground waters.
Improper disposal of municipal and industrial wastes in landfills and
land applications of sludges, industrial wastes, fertilizers, or pesticides
can lead to degradation of ground-water quality through leaching of these
materials and percolation of the contaminated leachate into shallow
aquifers. Aquifer contamination may be partially mitigated by natural
physical processes, but the time scale for flushing a conservative substance
from an aquifer is proportional to the hydraulic residence time of the
ground water, which can be hundreds of years in some aquifers. Furthermore,
the time required for the flushing of a hydrophobia contaminant that is
strongly sorbed by the aquifer solids may be much longer than the residence
time. Clearly, physical processes for aquifer self-purification can be very
slow. However, natural chemical and biological processes can also contri-
bute to aquifer self-purlfloat ion.
The rates of many of these self-purification processes depend on the
oxidation-reduction (redox) atatus of the ground-water/aqulfer system. The
redox status of the ground water can control the chemical speciation of many
elements and can also Influence microblal ecology and metabolism. Fo-
example, In the pH range of most ground waters, ferric hydroxide is much
less soluble than ferrous hydroxide. Also, in oxic systems containing
hydrous metal oxides, sorption processes limit the solubilities of many
trace elements [Jenne 1968]. For example, arsenate, which is the thermo-
dynamically favored form of arsenio In oxic waters is more strongly adsorbed
on metal oxides than arsenite, which is thermodynamically favored in anoxlc
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waters [Ferguson and Gavls 1972]. Arsenite is, thus, more mobile than
arsenate In aquifer systems. Arsenite is also more toxic than arsenate
[Lemmo et al. 1983]. Thus, in aquifers in which arsenic-containing minerals
are present, redox conditions are very Important in determining water
quality. In very anoxlc waters, precipitation of sulfides controls the
solubilities of many elements. The solubilities of such toxic metals as
cadmium and lead are likely to be highest in mildly anoxlc environments that
contain neither sulfide nor metal oxides, which are found in many shallow
aquifers.
Redox conditions also Influence mlcrobial speciation and metabolic
rates. Different classes of bacteria use different elements and compounds
as electron acceptors in respiration [Stumm and Morgan 1981, Stanier et al.
1979], If dissolved oxygen is present in an aquifer, then aerobic organisms
predominate. If aerobic respiration depletes the dissolved oxygen in a
system that is closed to the atmosphere, such as the deeper parts of an
aquifer, then a succession of microbial populations utilizing electron
acceptors with decreasing redox potentials may be expected. A typical ^micro-
bial succession may Include aerobes, denitrifiers, fermentors, sulfate
reducers, and methanogens. Such microbial successions have been invoked to
explain the chemical evolution of ground waters in several aquifers [Champ
et al. 1978]. The availability of preferred electron acceptors will affect
nicrobial speciation and the rates of biotransformation of toxic substances.
For example, the rate of degradation of DDT Is much greater in anoxic fresh-
water sediment systems than in oxic systems [Cambrell et al. 198H and refer-
ences cited therein].
In view of the importance of the redox status of ground water and other
natural waters it is desirable to have a convenient, reliable indicator of
redox status. The potential of a platinum electrode, or Eh, is often used
as a geochemical redox Indicator [Carrels and Christ 1965]. This is because
the Eh of a water sample can be readily measured and in veil defined
synthetic solutions the Eh provides a quantitative indicator of redox speci-
ation [Laitinen 1960]. Also, in certain aquatic environments the Eh can be
quantitatively related to chemical speciation, e.g. waters of low pH and
high Fe concentrations [Nordstrom et al. 1979] and certain anoxic sediments
[Emerson 1976]. However, using the Eh as a redox potential for an entire
aquatic system is only meaningful in systems that are in chemical equili-
brium and that contain electroactive solutes at approximately mi111molar
concentrations [Stumm and Morgan 1981], Most ground waters are not in
equilibrium [Lindberg and Runnels 1981]. Furthermore, the Eh of natural
waters are often mixed potentials determined by two or more redox couples,
rather than one dominant couple [Stumm and Morgan 1981]. Thus, it is usually
not possible to quantitatively relate chemical speciation and Eh in natural
waters.
In spite of its limitations, the Eh can be a useful qualitative indi-
cator of the redox status. Consistently measured or calculated Eh values
can Indicate relative redox levels in a single system. Such systems may
Include successive depths in a sediment or zones in flow systems with simi-
lar concentrations of major electrolyte ions and electroactive minor ions.
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On the other hand, comparison of the Eh values of very different waters,
e.g. well-poised anoxic ground waters and poorly-poised oxic ground waters,
is probably not meaningful. The most reliable characterization of the redox
level of a natural water is a complete chemical analysis, including all
redox-active species [Hostotler 1981*]. This is the approach which has been
taken in this project.
We have been collecting ground water samples from a pristine sand and
gravel aquifer in the Sand Ridge State Forest near Havana, Illinois. Using
the materials and techniques that are the least likely to disturb the
chemistry of the ground waters, we have measured Eh, redox speciation, and
complementary chemical parameters that enable calculation of redox condi-
tions in the aquifer. We will report our results and our interpretations,
including calculation of redox potentials from analytical data, redox
processes causing changes in dissolved oxygen concentrations, the presence
of hydrogen peroxide, and physical mixing in the aquifer.
CONCLUSIONS
The relative redox status of ground water at the Sand Ridge site is
related to the dissolved oxygen concentration as shown by Eh measurements
and chemical speciation calculations. The Eh also qualitatively indicates
the relative redox status of the ground waters, i.e. the lower the DO the
lower the Eh. In the deepest waters sampled, the Eh is quantitatively
related to Fe speciation. However, for other redox couples and at other
depths, there is no quantitative relationship between redox speciation and
Eh. The aquifer-ground water system is not in redox equilibrium at any
depth sampled because calculated redox potentials for several couples differ
by up to hundreds of millivolts, corresponding to differences in activity
quotients of orders of magnitude. The concentration profile of DO is fairly
constant and Is consistent with mixing of layers of oxic and anoxic water.
Modeling the DO profile may help understand mixing in the aquifer.
RECOMMENDATIONS
The field studies should be expanded to include very anoxic ground
waters in a similar hydrologic environment, preferably in the same aquifer.
Solid oxldants and reductants in the aquifer can be studied to determine
their role in redox processes. Geochemical extractions can give an estimate
of the aquifer oxidizing or reducing capacity. Redox titrations of aquifer
sediments can estimate the redox buffer capacity of the aquifer system.
Respirometry experiments may estimate kinetics of aquifer redox processes.
Characterization of the organic matter In ground water Is essential to
understanding aquifer redox processes In both oxic and anoxic waters. This
characterization should Include complexation of electroactive metals, which
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may influence measured Eh values, and molecular weight and functional group
determinations, which can indicate potential substrates for raicroblal
respiration.
MATERIALS AND METHODS
WELL DRILLING AND INSTALLATION
The monitoring wells for ground water sampling were installed in the
Sand Ridge State Forest near Havana, Illinois in the Havana Lowlands region.
The aquifer being studied is a water table aquifer composed of coarse sand
and gravel. The geology and hydrogeology of the area have been described by
Walker et al. [1965]. Hydrologic investigations have been conducted in the
area by the Illinois State Water Survey [Naymik and Sievors 1983, 1985].
Three wells were drilled in October, 1984 to nominal depths of 35, 50,
and 65 feet below grade. One additional well was drilled in September, 1985
to a depth of 104 feet. The wells were drilled with a hollow-stem auger. The
auger flights, well casing sections, and well screens were steam cleaned to
minimize introduction of foreign matter into the boreholes. During the
drilling of the 65-foot and 104-foot wells, split spoon samples of the
aquifer material, i.e. coarse sand or sand and gravel, were taken for deter-
mination of Fe and Mn oxide contents.
For each well the casing and screen were lowered by han-' through the
hollow auger stem. When the auger was withdrawn, the sand collapsed around
the screen and casing up to the water table, approximately 27 feet below
grade. Bentonite pellets were poured into the borehole to form a two-foot
barrier to rapid percolation. The hole was then backfilled with drill
cuttings to two to three feet below grade. Finally, the hole was filled
with expanding cement and a steel locking well protector was inserted and
allowed to set in the cement. Details of construction of the wells are
shown in Figure 1 .
The 50-, 65-, and 101-foot wells were developed by air lifting. An air
hose was lowered to the bottom of the well and compressed air was used to
force the water out of the well until the expelled water was clear. The
water level in the 35-foot well was too low to allow sufficient submergence
for development by air lifting. This well was developed by swabbing.
A Teflon positive-displacement bladder pump connected to Teflon tubing
was installed in each well. Before installation the pumps and tubing were
detergent and acid-cleaned followed by thorough rinsing with deionized
water.
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WATER SAMPLING
For pH, Eh, specific conductance, and temperature measurement ground
water was pumped from the well and through a flow-cell that held the
electrodes and probes [Garske and Schock In press]. The cell was designed so
that there would be no air contact with the ground water and so that the
axes of the pH and Eh electrodes were aligned parallel with, rather than
perpendicular to, the direction of flow. This cell design minimizes
streaming potential problems in pH and Eh measurements [Gray 1985] as well
as preventing accumulation of bubbles in the cell. The pH electrode waa
calibrated at the temperature of the ground water using two buffers.
The Eh was calculated from redox electrode (Orion 977800) potentials
using an empirical equation for reference electrode potential as a function
of temperature [Garske, E. E. unpublished data, Illinois State Water Survey,
J98UJ. The response of the two Eh electrodes waa also checked using a redox
buffer at the temperature of the ground water. The Eh values calculated
using the empirical equation differed from Eh values calculated relative to
the buffer by +30 to +80 mV.
Beginning in August, 1985 a polarographic dissolved oxygen (DO) probe
(Orion 970800) mounted in' a flow-through cell was Installed immediately
downstream from the flow-cell. The DO probe waa calibrated according to the
manufacturer's Instructions. Conductivity, pH, Eh, and temperature were
monitored in the flowing water. When all parameters reached stable values
(i.e. less than 0.05 pH unite, 10 mV, or 10 ohms"1 change in successive
casing volumes) the well was considered to be completely flushed. That is,
the water that was being pumped from that point on was considered to be
representative of the aquifer water. These stable values were recorded.
After well flushing, water samples were collected. Unfiltered water
samples were collected for organic carbon, ammonia, hydrogen peroxide,
hydrogen sulfide, and dissolved oxygen determinations. Organic carbon
samples were collected In precombusted glass vials. Samples for ammonia were
preserved with f^SOij. Samples for hydrogen sulfide were preserved with zinc
acetate and sodium hydroxide. Samples for dissolved oxygen and hydrogen
peroxide were analyzed immediately after collection. Oxygen diffusing
through the sampling tubing probably did not appreciably contaminate any of
our samples from the 35-, 50-, or 65-foot wells, but may have contaminated
the samples from the 101-foot well (see Appendix A).
After collection of the unfiltered samples, the sampling tubing was
connected to a 90 mm diameter In-line filter holder (Mllllpore) containing a
membrane of nominal pore size 0.1 ym (Nuclepore 1^1705). (Note: the hose
barbs supplied with the filter holder were replaced "by pressure-tight tubing
fittings.) The in-line filter holder was used to prevent air contact during
filtration, which can cause contamination of ground water samples by atmos-
pheric oxygen accompanied by oxidation of ferrous iron [Stolzenberg and
Nichols 1985]. Filter membranes of nominal pore size 0.1 urn were used
rather than the more commonly used 0.45 inn filters because'0.1 jim filters
are more effective for the removal of fine particulate metala from water
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samples [Kennedy et al. 1974, Laxen and Chandler 1982]. A new filter
membrane was used for each well. At least 500 mL of water was allowed to
flow through a new membrane to remove soluble contaminants and to equili-
brate the filter surfaces with dissolved trace metals, thus minimizing
adsorptive losses [Gardner and Hunt 1981]. After the initial washing, a 500
mL filtered sample was collected for alkalinity determinations. Filtration
of the alkalinity sample provided further washing of the filter. Subsequent
filtered samples were collected for major cations, anions and silica, and
trace metals. Samples for major cations and for trace metals were collected
in acid-cleaned bottles and were preserved with 1 percent (by volume) HNOj
and 0.1 percent HC1, respectively.
Field blanks were collected using the apparatus shown in Figure 2.
Deionized water was forced through the filtering apparatus and collected by
the same procedure as the water samples. Thus, the field blanks and samples
were treated identically.
Water samples were stored on ice immediately after collection. Upon
arrival at the laboratory the samples were refrigerated. Nonacidified
samples were stored for less than 24 hours before being analyzed.
ANALYTICAL METHODS
Determinations of the unstable solutes dissolved oxygen, alkalinity,
and hydrogen peroxide were performed in the field. Dissolved oxygen was
determined using the azide modification of the Kinkier method [Rand et al.
1975]. Alkalinity was determined by potentiometric titration using Gran's
method to locate the equivalence point [Stumm and Morgan 1981].
Manual colorimetric methods were used to determine ammonia [Standard
Methods 1975] and iron [Stookey 1970]. Both solutes were below detectable
levels in early samples. For the 9/19/85 and 10/17/85 samples from the
101-foot well, Fe and Mn were determined by atomic absorption spectrophoto-
metry. After 12/13/84, samples were screened in the field for NHj using a
field colorimetric method (Chemetrics, Inc. Calverton, VA). If the NHj
concentration was less than 0.1 mg L""1, the detection limit of the
Chemetrics method, the sample was not analyzed for NHj.
Automated adaptations of standard colorimetric methods were used to
determine orthophosphate, dissolved silica, sulfate, nitrite, and nitrate.
Chloride was determined by automated potentiometrio titration. Volatile and
nonvolatile organic carbon fractions were determined by wet oxidation and
infrared C02 detection [Barcelona 1981].
Manganese determinations were attempted using two electroanalytioal
methods. From November, 1981 until March, 1985 cathodic stripping voltam-
metry [Huber and Lemmert 1966, Hrabankova et al. 1969] using a
wax-impregnated graphite electrode. Manganese concentrations in ground
water samples were not significantly different from blank Mn concentrations.
From 4/23/85 to 6/20/85, anodic stripping voltammetry using a hanging
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mercury drop electrode [O'Halloran 1982] wa3 used. However, manganese was
not detected In any sample. After the 6/20/85 sampling trip, Mn determina-
tions were discontinued because the concentrations were too low for quanti-
tation.
Hydrogen Peroxide
The fluorlmetric scopoletin/horseradish peroxidase method [Andrae 1955]
incorporating published modifications [Perschke and Broda 1961 , Van Baalen
and Marler 1966, Cooper and Zika 1981] and further optimization developed
during this project, was used for' hydrogen peroxide determinations. Ths
method has been used to measure 2 nM HgOg in rain water [Zika et al. 1982],
and recently, H2C>2 in ground and surface waters [Cooper and Zika 1983]: The
method has several advantages over other methods for ^02: 1. It is
extremely sensitive, being based on a compound which has detectable
fluorescence in concentrations lower than 1 nM; 2. It is rapid, with
immediate a; nplete reaction due to the large turnover rate of the enzyme
catalyst; 3. ^t is selective, since the fluorophore is stable unless
oxidized by activated enzyme and the enzyme itself is highly specific for
and activated only by peroxides. (The contribution of organic peroxides can
be determined separately if necessary.); 1. The stoichiometry has been
established — one molecule of fluorophore is oxidized for each molecule of
H2C>2; and 5. It is adaptable to field determinations, eliminating problems
of sample instability frequently encountered in determinations of very low
levels of H202-
In the analysis, hydrogen peroxide present in the sample stoichiometri-
cally oxidizes scopoletin, a fluorescent lactone. The reaction is catalyzed
by the peroxidase enzyme. The fluorescence of the buffered mixture of sample
and reagent la measured before and after the addition of the enzyme, the
decrease or quenching of fluorescence being proportional to the amount of
hydrogen peroxide present. The fluorescence quenching is related to the
concentration of hydrogen peroxide by the method of standard additions.
Linear regression on a plot of fluorescence vs. moles added peroxide gives
the fluorescence response per mole of hydrogen peroxide reacted. This
response factor is used to calculate the amount of hydrogen peroxide in the
ground water sample. Ground water samples taken at Sand Ridge were
analyzed in triplicate within minutes of collection. At the low observed
concentrations, i.e. near the detection limit, It Is important to minimize
the imprecision of the concentration determinations and the potential for
In order v.o correct for any HgC^ contribution from the reagents, we
initially anaij^ed blanks, substituting de ionized, distilled, and freshly
redistilled water for the sample. However, hydrogen peroxide concentrations
in the reagent blanks were equal to or up to three times greater than those
of the samples for all sampling runs except 1/17, 3/19, 5/15, and 10/17/85.
We suspected that the distilled water contained 'HgOg and tested this
hypothesis by varying th« oonointralions of soopolatin una pwoxiaas* and by
adding the peroxidase before the scopoletin. Varying the reagent concentra-
tions had no effeot on the blanks and the H202 signal in the blanks was
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eliminated by addition of the enzyme prior to the scopoletin. Therefore,
the measured quenching in the reagent blanks was due to ^02 present in the
distilled water.
Hydrogen peroxide contamination of distilled water may be the major
problem with obtaining reliable reagent blanks. The stoichiometry of the
reaction in distilled water blanks has been found to be 0.7 to 0.9 moles of
scopoletin per mole of hydrogen peroxide, which is close to the accepted
value of 1.0 [Perschke and Broda 1961]. This supports the hypothesis of
^2^2 contamination of the distilled water and suggests that side reactions
are not responsible for quenching.
Blanks have not been addressed in the literature on this method.
However, a number of authors have suggested that f^Og can be produced in
deionized distilled water as a result of microbial activity [Lazrus et al.
1985], photochemical reactions [Zika 1984], or by sparging with air [Zika et
al. 1982]. Perschke and Broda [1961] suggested successive distillation of
blank and reagent water from KMnOjj,' ' AgNOg, and finally from a quartz still
to achieve sufficient water purity. However, none of the previous workers
reported any problems with background 1^2 concentrations in their experi-
ments.
We have further investigated the origin of the f^Og contamination in
the distilled water. We did not observe increases in the levels of H202 In
the water during storage over those in freshly distilled water. Also,
sparging the water with oxygen-free nitrogen for several hours did not
reduce the levels of #2^2 either. The most likely source of the ^02
contaminant is in the distillation process or is carried over from the
original amount in the deionized water.
Only one set of measurements of ^03 in distilled water using the
scopoletin/peroxidase method has been reported [Perschke and Broda 1961]. In
this1 paper the water was redistilled from KHnOj|, AgN03, and finally, -from a
quartz still. Unfortunately, since the experiments were intended only to
establish the stoichiometry of the reaction, the background contribution of
H202 in the distilled water was not reported. It is likely that hydrogen
peroxide is a normal trace component of our distilled water, produced from
the water itself and varying with inputs of radiation and dissolved oxygen.
Therefore, while a large fraction of our H202 determinations do not satisfy
the conventional analytical criterion of low blanks, we feel that we have
positively identified 1^2 in ground waters from the Sand Ridge site.
RESULTS AND DISCUSSION
The results of chemical analyses of ground water samples from the moni-
toring wells in the Sand Ridge State Forest are presented in Appendix B.
Aqueous ammonia, sulfate, phosphate, and dissolved silica concentrations are
expressed as milligrams of solute per liter (mg L~1). For example, a
%
8
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sulfate concentration of 1 mg L~1 is equivalent to 0.0104 millimolar (mM).
For nitrate and nitrite the concentrations are expressed as mg N per liter,
i.e. a concentration of 1 mg NOj" per liter is equivalent to 0.0711 mM.
Alkalinlties are in units of milliequivalents per liter (meq L~^) and can be
converted to milligrams per liter as CaCOj by multiplying by 50.
DISSOLVED OXYGEN, TRACE METALS AND EH
The dissolved oxygen (DO) profile for 9/19/85 is shown in Figure 3. The
profile of the upper three wells was typical of all sampling runs. The
35-foot sample had the highest DO concentration, which was near the DO
saturation limit at the ground-water temperature. The 50-foot sample was
near DO saturation with about 1 mg L~1 less than in the 35-foot sample. The
65-foot sample had less than 1.5 mg L~* DO. The DO measured in the deepest
sample was approximately- 0.2 mg L~^. However, because of the presence of
Fe2* and Mn^* in the deep waters, (described below) the measured DO was
probably an artifact of sampling. The suspected DO contamination nay have
resulted from gas diffusion through the sampling tubing (Appendix A). Thus,
0.2 mg L~1 may actually be an upper limit to the DO concentration in the
deep ground waters. The gradient in DO may be caused by mixing of recharge
water saturated in DO with deep ground water that is anoxic and contains
solutes, such as Fe^*, that consume DO. The time series graph of DO (Figure
4) shows fluctuations of up to 1.5 mg L""1 in one month at all depths. This
is greater than the uncertainty t)f +/- 0.2 mg L~1 in the Kinkier DO determi-
nations, so the observed fluctuations were not an artifact of the analyses.
There were no temporal trends in DO at any depth.
The Eh in the aquifer is related to the DO as shown by comparing the
profiles of DO and Eh (Figures 3 and 5, respectively). The Eh values are
high in oxic waters from shallow and Intermediate depths and low in the
nearly anoxic deep waters. The relative values of Eh measured in the 35-,
50-, and 65-foot wells were typical of all sampling trips with the highest
Eh measured in the waters from 35 feet, the lowest Eh in the 65-foot waters,
and a range of less than 50 mV. The tine series graph of Eh values (Figure
6) shows that, with the exception of 2/19, 3/19, and 9/19/95 the Eh in the
three shallow wells varied between +330 and >130 mV. The spread in Eh
values for the 35-, 50-, and 65-foot wells was nearly constant for all
sampling trips and agreement between duplicate readings was usually good.
Iron and Mn concentrations were below the detection limits of the
ferrozlne colorimetric and differential pulse anodic stripping voltammetrlo
m^'.nods, approximately 2 and 1 yg L""1, respectively, In all samples from the
35-, 50-, and 65-foot wells, in the samples from the 101-foot well, Fe and
Mn concentrations wort tpproxlmttly 0,5 ind 0,2 ng I"1, roapootivily,
Assuming that most of the dissolved Fe is ferrous; the partial pressure of
oxygen In equilibrium with this Fe24 concentration at the pH of the 101-foot
well is approximately 10**57 atmospheres, more than 50 orders of magnitude
less than the partial pressure calculated from the measured DO concentra-
tion. Using the rate law for the oxidation of ferrous iron by dissolved
oxygen [Stumra and Lee 1961, Singer and Stumm 1970, Morgan and Stumm 1961],
-------
the half-life of ferrous iron In the aquifer should be less than two hours,
which is certainly much shorter than the hydraulic residence time. These
discrepancies were probably the result of oxygen contamination during
sampling.
HYDROGEM PEROXIDE
The time aeries graph of hydrogen peroxide is shown in Figure 7. The
highest concentrations were measured on 1/17/85. The 2/19/85 concentrations
were somewhat lower than those observed on 1/17. There was no consistent
profile for
Hydrogen peroxide is an Important intermediate in the reduction of
oxygen in natural waters, a reaction which may be written as the product of
two two-electron reductions as shown in equations 2-5 [Breck 197*0
02 + 2 H* + 2 e~ --> H202 (2)
H202 - H* * H02~ (3)
H202 •«• 2 H* + 2 e- --> 2 H20 CO
H02~ * H20 * 2 e- --> 3 OH" (5)
where arrows indicate reactions proceeding in one direction and an equals
sign Indicates equilibrium. A steady state concentration may develop as the
result of balancing production and consumption reactions.
The 02/H202 couple may be a significant participant In the redox
chemistry of certain natural waters, but its Influence has not been studied.
In principle, the oxidizing power of dissolved oxygen can be controlled by
the kinetics of Its reduction; if the rate of reduction of H202 is slower
than the rate of its formation, then the potential is effectively that of
reaction 4 and 02 becomes a weaker oxldant than If it were directly reduced
to H20 rstumo and Morgan 1981 J. The reactivity of H202 Is limited by the
stability of th« O-O bond, which has half the strength of a single oovalent
bond. Although the dlsproportlonatlon of H202 to H20 and 02 is favored
thennodynamically (delta G - -234 KJ mole"*1), It is klnetloally slow In the
absence of trace metal or enzyme catalysts [Hoffmann 1984]. Thus, the low
Fe and Mn concentrations In the oxlc shallow ground waters at Sand Ridge may
contribute to the persistence of H202.
The contribution of the 02/H202 couple to the redox potential measured
in natural waters may be important [Breck 1974]* but has not been assessed.
In the Sand Ridge ground waters, HgO^ concentrations may be comparable to
those of other electroactive solutes. For example, dissolved Fe concentra-
tions were below the detection limit of approximately 2 yg L~1 , or 36 nM,
which is similar to H202 concentrations measured on some sampling trips;
There is now evidence that H202 is formed, and accumulates, in the
photooxldatlon of organic compounds in surface and ground waters [Cooper and
10
-------
Zika 1983]. However, the presence of H202 in untreated ground water has not
previously been reported. The superoxide radical anion has been implicated
as a precursor of H202 [Cooper and Zika 1983]. It may be formed by many
reaction pathways in natural waters, including the reduction of dissolved
oxygen by trace metals [Zika 1981]. Hydrogen peroxide is formed microbially
as a by-product of the destruction of the toxic superoxide radical anlon
[Stanier et al. 1979]. Hydrogen peroxide might accumulate in water due to
its slow kinetics of decomposition or the presence of tolerant organisms
lacking peroxidase or catalase enzymes for its destruction. For example,
bottom waters of the Volga River contain hydrogen peroxide concentrations
1.5 to 2 times higher than in surface waters, due to biochemical processes
[Sinel'nikov and Liberman 1971]. Since ^2^2 concentrations in precipitation
are frequently orders of magnitude greater than in ground water [Zika
et al. 1982] due to atmospheric photoproduction, rain water may also be a
source of peroxide in ground water (i.e. relict H2(>2 from recharge).
Van Baalen and Marler [1966] first measured HgOg in unfiltered sea
water, suggesting that It might be a "significant ecological variable".
Concentrations from 15 to 200 nanomolar were found in the surface ocean.
From samples taken at several sites they concluded that ^2Q2 was a
ubiquitous solute in sea water In the sampling area. No evidence was avail-
able as to the source of the oxidant, but atmospheric photoproduction and
precipitation, photocatalysls by pigments In the open water, and microbial
activity were suggested. Kok [1980] measured H202 in rain water, finding 5
to 17 micromolar H202 and a dependence of the concentration on "photochemi-
cal activity prior to the rainfall". Zika et al. [1982] found variations in
rain water from south Florida and Bahama Islands from 11 to 75 micromolar.
Draper and Crosby [1983] reported levels from less than the minimum
detectable for their assay, 2 yM, to 30 pM in irradiated, highly eutrophic
surface waters. The photoproduction of H202 from naturally occurring
organlcs, including humic compounds, was observed. Cooper and Zika [1983]
exposed unfiltered surface water and ground water to sunlight and determined
that H202 photoproduction was a function of the total organic carbon content
of the water, specifically the concentration of humics. Measurements of H202
in eight ground waters prior to sunlight Irradiation were below unreported
minimum detectable concentrations. They were probably less than 5 nM [Zika
et al. 1985]. The identification of superoxide as the H202 precursor
suggested to the authors that other transient reactive species, such as
hydroxyl radical and "organic radicals and metastable intermediates of many
transition metals and nonmetallic elements" might be involved in the redox
chemistry of natural waters. It was concluded that the levels of H202
measured in irradiated waters could affect the poise, or redox buffering
Intensity, of the systems. Studies of H202 photoproduction In several ground
waters and surface waters Indicated that there was little production of H202
in waters with low UV absorbance at wavelengths above 250 nm, corresponding
to low concentrations of humio substances [Cooper and Zika 1981], Despite
the wide-ranging sampling and detailed analysis of H202 photoaccumulation
rates, the only measurements of H^ In ground waters prior to irradiation
were below detection limits [Cooper and Zika 1983].
11
-------
OTHER SOLUTES
The concentration profiles of Ca, Mg, alkalinity, and specific
conductance (Figure 8) show a sharp decrease between 35 feet and 50 feet
with a smaller decrease between 50 and 65 feet. These profiles were consis-
tent for all sampling runs. There was a small decrease between 65 and 101
feet In alkalinitles and concentrations of Ca and Kg. Th9 specific conduc-
tance profile is consistent with the profiles of the major ions. The time
series graphs of Ca, Mg and alkalinity are shown In Figures 9-11.
Alkalinlties and Ca, and Mg concentrations tend to Increase with time at
depths of 35 and 65 feet and are fairly constant at 50 feet. There are
concentration fluctuations of more than 0.5 mg L~1 month'1 at all depths.
The concentration profiles of Na, K, sulfate and chloride (Figure 12)
for were typical of all observed profiles. There was a decrease between 35
and 50 feet for all four solutes. Between 50 and 65 feet concentrations of
Na, K, and chloride concentrations changed by small amounts, either
Increasing or decreasing, while sulfate increased. Between 65 and 101 feet
Na and K increased slightly while chloride and sulfate changed very little.
Sodium and sulfate concentrations increased with time in the 35-foot
samples, similarly to Ca, Mg, and alkalinity, while K and chloride showed no
trend with time (Figures 13-16).
The pH profile (Figure 17) was typical of all sampling runs with pH
increasing from 35 to 65 feet. There was a decrease in pH from 65 to 104
feet. The pH of 35-, 50-, and 65-foot ground waters passed through a
maximum on 12/13/84 and varied only within 0.3 units for the rest of the
sampling runs.
The concentration profiles of nitrate and orthophosphate (Figure 18)
were representative of most sampling runs. Nitrate concentrations decreased
with depth, while phosphate concentrations increased. The decrease to a
nearly undetectable nitrate concentration in the 104-foot well is consistent
with the sharp decrease in DO and Eh between 65 and 104 feet. However,
reduced nitrogen species, nitrite or ammonia, were not detected. Nitrate
concentrations tended to decrease with time (Figure 19), while phosphate
concentrations showed no particular trend after a sharp drop between
The concentration profiles of DO, NO*", and , possibly, Q-POjj3" for
9/19/85 have two Inflection points. All three profiles are concave upward
from 35 to 65 feet and concave downward from 50 to 104 feet. Concentrations
of N03~ and DO decrease with depth, while o-POjj3* concentrations increase
with depth. Concentration profiles like these can be produced by
diffusion/dispersion. The conceptual model is that upgradient from the
monitoring wells two layers of ground water with different concentrations of
DO and nitrate are separated by a thin transitional region. Diffusion and
dispersion causes the transitional layer to broaden and produces the
observed profile. The analogous heat transfer problem, which is
mathematically identical, is that of two semi-infinite rods of the same
material but at different temperatures that are placed face to face.
12
-------
Solutions to this problem have been published [e.g. Churchill 19^1]. It may
be possible to estimate the average dispersion coefficient in the aquifer by
modeling the concentration profiles of DO, nitrate, and other solutes.
The time series graph of total organic carbon (TOG) in the 35-, 50-,
and 65-foot monitoring wells is shown in Figure 21. There are considerable
changes in TOC with time at all depths. Some of the TOC changes in consecu-
tive months were in the same direction, i.e. in one month the changes at 35,
50, and 65 feet were all increases or all decreases. In other months the
changes at different depths were in different directions. There did not
appear to be a consistent profile of TOC. Wood and Peraltis [1981] observed
that DO decreased and the partial pressure of CO2 increased with depth below
the active soil zone, probably due to microbial respiration. We compared
month-to-inonth changes in DO and TOC in an attempt to explain the observed
variations in DO, with significant correlation indicating control of DO by
microbial respiration. However, the relative uncertainty in TOC measure-
ments was too large to make meaningful correlations.
CALCULATING REDOX POTENTIALS FROM ANALYTICAL DATA
At equilibrium the potential of a redox couple is given by the Nernst
equation (equation 6)
*
E - E° + (RT In10 / nF) log(Ox/Red) (6)
where E is the potential, E° is the standard potential (a constant), R is
the gas constant, T is absolute temperature, n is the number of electrons
transferred, F is Faraday's constant, and Ox and Red are the product of
chemical activities of the oxidized and reduced aides of the atoionioroetric
equation, respectively. At 12eC, the temperature of Sand Ridge ground
waters, the value of RTlnlO/F is 56.6 mV. In a solution containing more
than one redox couple at equillbriumi all potentials calculated by equation
6 are equal. However, natural waters are rarely In redo* aqullibrlum iStuatt
and Morgan 1981]. Thus, the potentials calculated using equation 11 for a
natural water are strictly theoretical. In this section we describe how
potentials are calculated from analytical data and compare measured and
calculated potentials.
The redox couples considered were 02^2^2* Fe(III)/Fe(II), NOa~/NHij*,
Mn(III)/Mn(II) and SOjj2-/HS~. For each couple a range of potentials was
calculated based on ground water analyses in Tables 1-12 and these calcu-
lated potentials are compared to the observed range of Eh values.
The potential of the 02/^2^2 GOUPle is given by equation 7
E(mV) - 695 * 28.3 ( log(DO/54.1) - 2 pH - log[H2023 ) (7)
13
-------
where the standard potential at 12°C is 695 mV, square brackets indicate
molar concentrations, DO is dissolved oxygen in mg L~1 and 51.1 is the
factor that converts DO to oxygen partial pressure at 12°C. (Strictly
speaking, activities should be used in the Nernst equation, rather than
concentrations. However, 02 and ^2^2 are uncharged species, so their
activity coefficients were assumed to be 1.0. Because pH measures H+
activity, no further ionic strength correction is necessary.) The maximum
calculated 0^>/^2^2 potential was for water samples collected from the
35-foot well with a DO of 9 mg L~1 , and pH of 7.5. The minimum potential
was for 65-foot samples with DO of 3 mg L'1 and pH of 8. The DO concentra-
tions measured in the 101-foot well were not considered because it was felt
that the values were questionable. A. range of H^g concentrations of
0.1-10.0 nM was assumed for these calculations. The DO-H202 potentials
ranged from +130 to +525 mV, which overlaps the observed range of most of
the Eh measurements for the 35- to 65-foot ground waters, roughly +350 to
+1J50 mV.
For the NO^'/NHjj* couple only nitrate was detected. As a result, the
Nernst equation becomes an inequality (equation 8)
E(mV) > 881 + 7.08 ( log[N03~] -• 10 pH - log DL ) (8)
where the standard potential at 12°C Is 881 mV and DL is the detection limit
for ammonia nitrogen (0.05 mg L'1). (Because the activity coefficient for
singly charged ions at the Ionic strength of Sand Ridge ground waters was
0.92 as calculated by the Da vies approximation, the concentration of nitrate
was used in place of the activity.) The minimum nitrate concentration in
ground waters from 35 to 65 feet was 21 microoolar In a 65-foot sample. The
calculated lower bound for NOf-NHn* potentials was +321 mV, which la
consistent with the observed range of Eh values. If the KHjj* concentration
were less than 7 ng L~1 , then the calculated potentials would fall In the
observed Eh range. The NOj'-NHn* potential for the 101-foot ground waters
of +333 mV is far from the observed value of +110 aV.
Iron and manganese were below detection in all samples collected from
the 35-, 50-, and 65-foot wells. However, hydrous oxides of Fe and Mn were
found In aquifer sand samples, Thus, lower bounds for apparent potentials
of couples involving these metals can be calculated assuming equilibrium
with the oxides. For the couple Fe(OH)3/Fe2* the Nernst equation becomes the
inequality given by equation 9
E(mV) > 756 + 56.6 ( log K - 3 pH - log DL ) (9)
where a detection limit of 1 vg L~1 is assumed for Fe determination by the
ferrozlne method, all soluble Fe Is assumed to be Fe2*, and K is the equili-
brium constant for dissolution of Fe(OH)3. The value of K at 12°C Is 1.88
11
-------
[Langmuir 1969]. (Note that the only variable In equation 9 is pH.) The
lower bound for the Fe(OH)3/Fe2* potential was estimated to be + 155 raV,
which is not close to the observed Eh values, but Is at least consistent. If
the Fe2* concentrations were less than 20 ng L"1, and Fe3* levels were in
equilibrium with Fe(OH)j as assumed, then the calculated Fe(OH)3-Fe2*
potential would fall in the observed Eh range.
Iron was detected in the 10il-foot ground waters. It was assumed that
all soluble Fe was in the ferrous form. The activity of Fe2* was calculated
assuming equilibrium with sulfate and hydroxide complexes and using the
extended Debye-Huckel approximation to calculate the activity coefficient.
The activity of Fe2* was substituted for the detection limit in equation 9.
The calculated Fe(OH)3/Fe2* potential was +38 mV. An uncertainty in the
value of the solubility product of Fe(OH>3 of one logarithmic unit, a pH
reading that was high by 0.33 pH units, or a combination of smaller uncer-
tainties in these parameters would explain the disagreement between the
calculated and measured potentials. The solubility product determined at
25 °C was used for the potential calculation because the enthalpy change of
the reaction was not listed, which precluded a temperature correction. Thus,
the calculated Fe(OH)3/Fe2* potential agrees reasonably well with Eh. FeOOH
was also considered as the phase controlling Fe3*. Because the value of log
K for FeOOH is 0.98, the calculated potentials were lower than those calcu-
lated assuming FetOH)} control by approximately 200 mV. This disagreement
cannot be rationalized in terms of experimental uncertainties. Obviously, it
is important to know the form of hydrous Fe oxide%ppesent in the aquifer.
Hanganite, HnOOH, was assumed to be the form of hydrous Mn oxide
pr««& in «\i t^iVttr* For NnQQtt/Ntt** il\ tilt l^T XQ famCQQ\ frO\H\ 1491 * 56.6 ( -log DL - 3 pH ) (10)
As with the Fe calculations the only variable in this inequality is pH. The
lower bound for Kn potentials calculated according to equation 5 is +570 mV,
which is greater than all but two extreme Eh values. In order for the cal-
culated Mn potential to fall in the measured Eh range the Mn2* concentration
would have to be greater than 40 mg L~^, which would have been easily
detectable. The standard potential for Mn02/Mn2* is higher than that for
MnOOH/Mn2*, so the potentials calculated for the former couple disagreed
with measured potentials by an even greater amount. Manganese was detected
in the 104-foot waters. However, assuming the presence of MnOOH at 104
feet, the calculated Mn potential is +474 mV, which is much higher than the
observed value of +100 mV. Apparently the MnOOH-Mn2* couple is not at
equilibrium in the aquifer.
Measured and calculated potentials are compared in Figure 22. The best
agreement between calculated and measured potentials was for Fe In the
nearly anoxic 104-foot well. The potentials of the other couples differed
15
-------
from Eh by amounts that were not possible to explain by experimental uncer-
tainty. In the shallower, oxic wells, the range of potentials calculated
for the Q2~^2°2 couple overlapped the range of Eh values measured over one
year. However, for individual samples, the differences between calculated
and measured potentials correspond to differences in the ratio of Q£ to f^Og
concentrations of many orders of magnitude. Similar lack of agreement
between measured potentials and potentials calculated from analytical data
has been noted by Lindberg and Runnels [1984], in their examination of
published ground water data. In certain special situations measured Eh
values have been found to correspond to speciation of one or more elements.
These environments include waters of streams receiving acid mine drainage
and having low pH values and high Fe concentrations [Nordstrom et al. 1979],
anoxic ground waters having high Fe concentrations [Back and Barnes 1965],
and anoxic sediment interstitial waters having high sulfide concentrations
[Emerson 1976]. However, in oxic waters at near neutral pH the measured Eh
is usually far from that calculated from elemental speciation. The probable
reason for the lack of agreement between measured and calculated Eh is that
the system is not in chemical equilibrium.
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sea water. Nature, no. 5052, p. 951.
Walker, w. H., R. E. Bergstrom, and W. C. Walton. 1965. Preliminary report
on the ground-water resources of the Havana region In west-central Illinois.
Cooperative Ground-Water Report 3. State Water Survey, State Geological
Survey. Champaign, IL 61820.
Wood, W. W. and M. J. Peraitis. 1984. Origin and distribution of carbon
dioxide in the unsaturated zone of the southern high plains. Wat. Resour.
Res. 20:1193-1208.
Zlka, R. G. 198U. Short-lived oxidants in natural waters. Conference on
gas-liquid chemistry in natural waters. April, 1981. Brookhaven National
Laboratory. Extended abstract 52.
Zika, R., E. Saltzman, W. L. Chameldes, and D. D. Davis. 1982. Hydrogen
peroxide levels in rainwater collected In. south Florida and the Bahama
Islands. J. Geophys. Res. 87:5015-5017.
Zika, R. G., E. S. Saltzman, and W. J. Cooper. 1985. Hydrogen peroxide
concentrations in the Peru upwelling area. Mar. Chem. 17:265-275.
ZoBell, C. E. 1946.* Studies on redox potential of marine sediments. Bull.
Amer. Assoc. Petr. Geol. 30:177-513-
20
-------
APPENDIX A
OXYGEN DIFFUSION THROUGH SAMPLING TUBING
Oxygen is known to diffuse through polymeric materials, including
Teflon. (Teflon membranes are used in polarographic oxygen analyzers
[Hitchman 1978].) Thus, it is possible that our ground water samples were
contaminated by atmospheric oxygen. If this were true, then our dissolved
oxygen measurements would be biased and there could have been some oxidation
of ferrous iron before sample acidification. To estimate the extent of this
contamination we assumed that the gradient of oxygen fugacity across the
tubing wall is linear and equal to the difference between the fugacity
inside and outside the tube divided by the tubing thickness. Assuming
steady state in a completely mixed control volume of tubing and rearranging
terms yields equation A1
d f(02) / dx - (2 v r K A f) / (Q A r) (AT)
where f(02> is the oxygen fugacity, r is the distance from the center of the
tube,A f/A r is the gradient of D£ fugacity, Q is the flow rate, and K is
the mass transfer coefficient for 02 through Teflon. The dimensions of K
are (cm3 02 at STP) (cm~2 sec"1) (cm of Teflon) (cm of Hg)~1 and the magni-
tude is approximately 1 x 10~9 [Personal communication, William Buxton, Du
Pont Company, February 1985], Integration of equation B1 gives
z / z0 - exp(-2 K L / (KH Q)) (A2)
where z is the difference in Og fugacity (in atmospheres) between the atmos-
phere and the inside of the tube at a given point, L is the distance from
the beginning of the tube to the point of integration, and Zo is the value
of z at the beginning of th« tube, i.e. at L - 0. The difference in
fugacity is defined by equation A3
z - 0.2 - DO / KH (A3)
where, the approximate fugacity of 02 in air is 0.2 atmospheres, and KH is
the Henry's law constant for oxygen dissolution. For initially anoxic water
being pumped through 100 feet of tubing at a rate of 1000 mL min'1 the
measured DO would be 0.14 mg L"1, which Is similar to the observed DO
values. For waters having initial DO values of 3*5 and 8.5 mg L~1, the
calculated Increases in DO would be 0.10 and 0.03 ng L"1, respectively.
These increases are less than the uncertainty in the Kinkier titration.
21
-------
Using the rate law for ferrous Iron oxidation by dissolved oxygen
[Stumm and Morgan 1981] It Is possible to predict the rate of oxidation due
to 0? diffusion through sampling tubing. The resulting system of differen-
tial equations cannot be conveniently solved in closed form. We used a
fourth-order Runge-Kutta method [Kreyszig 1972] to solve the equations
numerically. The solutions for a number of Initial conditions (DO, Fe )
and pumping rates are presented In Table A1. The flow rates of 100 and 1000
mL rain"1 are typical of flow rates during ground water filtration and
sampling for DO measurements, respectively. The DO concentrations of 0,
3 5, and 8.5 mg IT1 are typical of ground waters from 10H, 65, and 35 feet
depths, respectively. If this model Is correct, then oxygen contamination
can seriously affect the determination of DO, Fe speciatlon, and, possibly,
Eh in mildly anoxic ground waters. We are planning experiments to test the
diffusion hypothesis.
TABLE A1. OXYGEN DIFFUSION THROUGH SAMPLING TUBING
Initial DO Initial Fe2* • • DO Fe2*
Concentration Concentration Flow Rate Increase Decrease
(mg L~1) (pM) (ml rain"1) (ml min"1) (wM)
0.0
0.0
0.0
0.0
3.5
3.5
8.5
8.5
0.00
1.00
0.00
1.00
0.00
0.00
0.00
0.00
100
100
1000
1000
100
1000
100
1000
1.36
1.36
. 0.11
0.1H
0 ,94
0.10
0.33
0.03
KM
0.85
—
0.99
—
—
VM
—
22
-------
APPENDIX B
TABULATION OP GROUND WATER ANALYSES
TABLE B1. ANALYSES OF GROUND WATER SAMPLES COLLECTED 11/13/81
Depth (feet)
Temperature (degrees C)
PH
Eh1a (mv)
Eh2a (mv)
Conductivity (yS/cra)
Alkalinity (meq/L)b
Ammonia
Nitrite
Nitrate
SulfaU
Chloride
Orthoptiosphate
Sulfide
Dissolved Silica
Purgeable Organic Carbon
Nonpurgeable Organic Carbon
Calcium
Magnesium
Sodium
Potassium
Iron (yg/L)
Manganese (\ig/L)
35
13.2
7.90
379
379
310
3.69
<0.05
<0.01
1.72
S8»9
2.38
0.11
<0.025
11.1
1.6
0.7
55.7
19.1
2,66
0.59
<0.1
<0.05
50
13.1
8.09
371
357
255
3.20
<0.05
<0.01
1.16
\9»7
1.88
0.16
<0.025
11.1
1.7
0.1
15.5
16.5
2.23
0.57
<0.1
<0.05
65
12.0
8.19
371
371
208
2.15
<0.05
<0.01
2.03
\9tl
1.96
0.17
<0.025
13.3
1.1
0.1
36.7
13.1
1.97
0.17
<0.1
<0.05
Notes: aE1, E2 are Eh values calculated according to
equation 1.
^Concentration units are mg/L except as noted.
23
-------
TABLE B2. ANALYSES OF GROUND WATER SAMPLES COLLECTED 12/13/83
Depth (feet)
Temperature (degrees C)
pH
Eh1 (mv)
Eh2 (mv)
Conductivity (pS/cm)
Alkalinity (meq/L)
Amn»nla-N
Nitrite
Nitrate
Sulfate
Chloride
Orthophosphate
Total Phosphate
Dissolved Silica
Nonpurgeable Organic Carbon
Purgeable Organic Carbon
Calcium
Magnesium
Sodium
Potassium
Iron (ug/1)
Manganese (pg/1)
Dissolved Oxygen
Hydrogen Peroxide (nM)
35
11.1
8.20
338
336
272
3.81
0.1 Up
<0.01
1.20
3*».8
1.95
<0.050
<0.050
15.2
0.7
2.0
56.0
19.6
2.8
0.8
<2.0
<1.0
9.30
6.9
50
10.3
8.U6
311
339
235
3.42
0.098
<0.01
1.42
19.7
1.33
<0.050
<0.050
15.1
0.7
1.6
47.6
17.2
2.4
0.8
<2.0
<1.0
8.00
2.8
65
9.7
8.66
338
335
180
2.45
0.071
<0.01
1.56
19.7
1.60
0.055
0.067
13.9
0.7
2.2
36.0.
12.8
2.0
0.4
<2.0
<1.0
4.21
6.5
-------
TABLE B3. ANALYSES OF GROUND WATER SAMPLES COLLECTED 1/17/85
Depth (feet)a
Temperature (degrees C)
PH
Eh1b (mv)
Conductivity (uS/cm)
Alkalinity (meq/L)
Amraonia-N
Nitrate
Sulfate
Chloride
Orthophosphate
Dissolved Silica
Nonpurgeable Organic Carbon
Purgeable Organic Carbon
Calcium
Magnesium
Sodium
Potassium
Iron (vg/1)
Manganese (wg/1)
Dissolved Oxygen
Hydrogen Peroxide (nM)
35
8.9
7.73
293
224
3.79
<0.05
1.16
30. 34
2.09
<0.05
15.2
0.5
1.7
56.4
19.6
2.8
0.8
<2.0
-------
TABLE B4. ANALYSES OF GROUND WATER SAMPLES COLLECTED 2/19/85
Depth (feet)
Temperature (degrees C)
PH
Eh1 (mV)
Eh2 (mV)
Specific conductance (uS/cm
Alkalinity (roeq/L)
Nitrate
Sulfate
Chloride
Dissolved Silica
Prugeable organic carbon
Nonpurgeable organic carbon
Calcium
Magnesium
Sodium
Potassium
Iron (pg/L)
Dissolved Oxygen
Hydrogen Peroxide (nM)
35
10.5
7.65
532
553
320
4.43
1.58
35.7
2.70
15.5
3.9
3.8
58.0
20.2
2.8
0.8
<2.0
10.2
23.0
50
10.2
7.96
539
562
250
3.64
1.43
17.9
-2.17
• 15.1
•1.1
1.3
44.8
16.7
2.0
0.4
<2.0
9.4
2.9
65
10.1
8.16
517
538
210
2.77
1.17
22.7
5.70
11.2
0.9
1.1
36.4
13.6
2.0
0.4
<2.0
5.3
7.5
26
-------
TABLE B5. ANALYSES OF GROUND WATER SAMPLES COLLECTED 3/19/85
Depth (feet)
Temperature (degrees C)
PH
Eh1 (mV)
Eh2 (mV)
Conductivity (yS/cm)
Alkalinity (meq/L)
Nitrate
Su.lfate
Chloride
Orthophosphate
Dissolved Silica
Purgeable Organic Carbon
Nonpurgeable Organic Carbon
Calcium
Magnesium
Sodium
Potassium
Iron (yg/1)
Manganese (yg/1)
Dissolved Oxygen
Hydrogen Peroxide (nM)
35
13.0
7.57
386
108
358
1.10
0*81
12.1
2.28
0.0173
16.3
0.6
1.2
58.3
21.1
2.6
0.7
<2.0
<1 .0
9.6
<1.0
50
11.5
7.82
386
107
270
3.02
KOI
18.5
0.68
0.0300
15.7
0.1
<0.1
13. 1
16.5
2.0
0.6
<2.0
<1 .0
8.6
1.3
65
13.2
8.01
372
388
221
2.12
1*60
22.3
1.10
0.0668
11.5
0.5
0.1
36.1
13.8
1.9
0.6
<2.0
<1 .0
3.7
1.2
27
-------
TABLE B6. ANALYSES OF GROUND WATERS SAMPLES COLLECTED 4/23/85
Depth (feet)
Temperature (degrees C)
PH
Eh1 (mV)
Eh2 (mV)
Conductivity (,\iS/cm)
Alkalinity (meq/L)
Sulfate
Chloride
Orthophosphate
Dissolved Silica
Purgeable Organic Carbon
MonpuTffMbl* Organlo Carbon
Calolu*
Magn«9liun
Sodlua
Potassium
Iron (ug/D
Dissolved Oxygen
Hydrogen Peroxide (nM)
35
13.0
7.55
321
359
3*0
3.86
HO.O
2.33
0.023
15.2
<0.1
0.2
60. t
20.4
2.?6
0.66
<2.0
9.8
<1 .0
50
13.2
7.71
321
359
265
3.07
20.0
1.26
0.035
13.8
<0. 1
<0. 1
44.7
16.0
2. OS
0.62
<2.0
8.6
1.3
65
n.i
7 '.97
31*
342
225
2.H6
20.0
1.50
0.072
15.0
0.2
0.4
37.*
13.4
1.91
0.54
<2.0
1.5
<1 .0
-------
TABLE B7. ANALYSES OF GROUND WATER SAMPLES COLLECTED 5/15/85
,
Depth (feet)
Temperature (degrees C)
PH
Eh1 (mV)
Eh2 (mV)
Conductivity (uS/cm)
Alkalinity (meq/L)
Nitrate
Sulfate
Chloride
Orthophosphate
Dissolved Silica
Purgeable Organic Carbon
Nonpurgeable Organic Carbon
Calcium
Magnesium
Sodium
Potassium
Dissolved Oxygen
Hydrogen Peroxide (nM)
35
12.6
7.17
121
113
350
3.95
1.13
11.1
1.95
<0.010
15.1
0.2
0.9
60.2
20.1
2.71
0.61
9.5
8.5
50
13.0
7.76
112
131
270
3.13
1.09
19.2
0.82
<0.010
11.9
0.1
0.6
11.3
15.9
2.10
0.60
8.8
2.3
65
13.1
7.95
107
127
230
2.53
0.77
23.3
1.02
0.053
11.0
0.2
0.1
37.2
13.2
1.93
0.52
1.2
1.9
29
-------
TABLE B8. ANALYSES OF GROUND WATER SAMPLES COLLECTED 6/20/85
Depth (feet)
Temperature (degrees C)
pH
Eh1 (mV)
Eh2 (mV)
Conductivity (uS/cm)
Alkalinity (meq/L)
Nitrate
Sulfate
Chloride
Or thophosphate
Dissolved Silica
Purgeable Organic Carbon
Nonpurgeable Organic Carbon
Calcium
Magnesium
Sodium
Potassium
Dissolved Oxygen
35
11.9
7.10
129
111
119
1.18
1.11
16.0
2.36
0.018
*
15.1
1.7
<0.1
61.8
21.6
2.88
0.65
9.1
50
12.7
7.67
123
109
120
3.19
1.22
21.7
1.25
0.016
11.9
0.2
0.1
15.2
16.7
2.19
0.61
7.9
65
13.0
7.90
318
361
70
2.66
0.15
27.6
1.11
0.088
11.0
0.1
<0.1
38.9
11.2
2.05
0.51
1.5
Note: Hydrogen peroxide analyses not performed.
30
-------
TABLE B9. AMALYSES OF GROUND WATER SAMPLES COLLECTED 7/23/85
Depth (feet)
Temperature (degrees C)
PH
Ehl (mv)
Eh2 (mv)
Conductivity (pS/cm)
Alkalinity (meq/L)
Nitrate
Sulfate
Chloride
Orthophosphate
Dissolved Silica
Purgeable Organic Carbon
Monpurgeable Organic Carbon
Calcium
Magnesium
Sodium
Potassium
Dissolved Oxygen
Hydrogen Peroxide (nM)
35
11.2
7.56
385
375
355
3.69
1.18
42.3
2.11
0.019
15.7.
<0.1
0.8
65.0
22.4
2.96
0.66
10.3
<1.0
50
11.9
7.83
379
367
258
2.97
0.97
23.7
2.18
0.037
15.2
<0.1
<0.1
43.7
16.1
2.14
0.60
9.4
5.8
65
12.2
8.04
371
357
237
2.59
0.56
31.6
1.41
0.028
14.2
<0.1
0.4
39.9
14.6
2.10
0.56
4.5
<1.0
-------
TABLE B10. ANALYSES OF GROUND WATER SAMPLES COLLECTED 8/22/85
,
Depth (feet)
Temperature (degrees C)
PH
Eh1 (mV)
Eh2 (mV)
Conductivity (uS/cm)
Alkalinity (raeq/L)
Nitrate
Sulfate
Chloride
Orthophosphate
Dissolved Silica
Purgeable Organic Carbon
Nonpurgeable Organic Carbon
Calcium
Magnesium
Sodium
Potassium
Dissolved Oxygen
35
11.9
7.57
433
435
317
4.37
1.07
44.5
2.31
0.0100
15.0
0.7 .
0.7
63.2
22.6
2.80
0.60
9.1
50
12.8
7.76
421
420
228
3.13
0.89
20.5
1.15
0.0167
11.5
0.6
<0.1
42.5
16.2
2.00
0.50
8.0
65
13.1
8.13
381
376
228
2.74
0.43
28.6
1.49
0.0670
13.9
0.7
0.3
41.0
15.4
2.00
0.50
4.2
Note: Hydrogen peroxide analyses not performed.
32
-------
TABLE B11. ANALYSES OF GROUND WATER SAMPLES COLLECTED 9/19/85
,
Depth (feet)
Temperature (degrees C)
PH
Eh1 (mV)
Eh2 (mV)
Conductivity (pS/cm)
Alkalinity (meq/L)
Nitrate
Sulfate
Chloride
Orthophosphate
Dissolved Silica
Nonpurgeable Organic Carbon
Calcium
Magnesium
Sodium
Potassium
Iron
Manganese
Dissolved Oxygen
Hydrogen Peroxide (nM)
35
12.5
7.60
568
555
325
1.16
1.06
43.8
2.45
0.004
15.7
0.7
65.6
22.7
3.00
0.66
<0.08
<0.03
8.5
4.5
50
13. 4
7.91
557
546
234
.2.98
0.80
19.9
1.14
d.021
15.1
0.2
42.3
16.0
2.04
0.58
<0.08
<0.03
7.3
1.7
65
13.8
8.08
524
504
230
2.73
0.34
29.4
1.42
0.051
14.3
0.1
41.9
15.6
2.13
0.54
<0.08
<0.03
3.5
1.7
104
15.9
7.77
114
116
223
2.39
0.04
28.6
1.30
0.101
16.4
0.4
38.6
12.3
3.06
0.72
0.49
0.18
0.2
2.3
33
-------
TABLE B12. ANALYSES OF GROUND WATER SAMPLES COLLECTED 10/17/85
Depth (feet)
Temperature (degrees C)
PH
Ehl (mv)
Eh2 (mv)
Conductivity (mS/cm)
Alkalinity (meq/L)
Nitrate
Sulfate
Chloride
Orthophosphate
Dissolved Silica
Calcium
Magnesium
Sodium
Potassium
Iron
Manganese
Dissolved Oxygen
Hydrogen Peroxide (nM)
35
12.0
7.76
418
394
345
4.52
0.86
46.0
2.58
0.04
15.4
68.2
23.9
3.08
0.68
<0.08
<0.03
9.7
27.5
50
11 .8
8.04
420
396
235
3.00
0.77
20.9
1.18
0.04
14.6
42.8
16.1
2.03
0.58
<0.08
<0.03
8.5
a
65
12.1
8.21
401
381
236
2.86
0.42
29.6
1.50
0.07
14.0
43.2
16.1
2.14
0.56
<0.08
<0.03
3.1
22.0
104
12.4
8.00
109
101
230
2.59
0.03
27.8
1.46
0.11
15.8
39.0
12.5
3.03
0.72
0.46
0.16
0.7
a
Note: aHydrogen peroxide only determined in 35- and 65-foot
samples.
-------
WELL DOM WELL DOM
8 Inch dlanwttr
WELL006S WELL 00105
20
m
Z
t
in
O 60
80
100 L-
X
Bmtonta
Plug
inttrval
77J
TIL
feu
Tabta
bonhote
ZZZ
Figure f. Construction of monitoring wells.
35
-------
Nitrogen
Pressure
Tubing Connector
(Swagelok)
To Filter Holder
Bored-through Fitting
(Swagelok)
-Bottle Containing
Deionized Water
• Pressure-tight
Plexiglas Vessel
Figxire 2. Apparatus for filter blanks.
36
-------
9
0
-10
-2O
-3O
-4O
-50
-60
-7O
-8O
-90
-100 H
-11O
4 6
DTs*olvftd Oxygan (mg/L)
i
8
Figure 3. Concentration profile of dissolved oxygen in
Havana lowlands aquifer, 9/19/85.
37
-------
c
9
I
o
1
o
M
J»
O
11
10
9 -
8 -
7 -
6 -
5 -
4 -
I
100
200
Time (days)
300
Figure 4. Tine aeries graph of diaaolved oxygen concentrations
in Havana lowlanda aquifer.
n 35, * 50. O 65 feet.
38
-------
0 -|
-10 -
-20 -
-3O -
-4O -
-50 -
-60 -
-70 -
-80 -
-90 -
-100 -
-11O -
1OO
200
30O 400
Eh (mV>
500
60O
Figure 5. Profile of Eh values In Havana lowlands aquifer, 9/19/85.
39
-------
v^
U
eoo
soo -
4OO -
300 -
2OO -
100 -
300
O 35
Tlm« (days)
50 o 65
A 104
Figure 6. Tine series graph of Eh values in Havana lowlands aquifer.
D35, * 50. O 65. A 101 feet.
-------
9
Q.
9
2
•o
20 -
10 -
100
200
Tim* (days)
300
4OO
Figure 7. Tlae series graph of hydrogen peroxide concentrations
in Havana lowlands aquifer.
D 35. * 50, O 65 feet.
JJ1
-------
0.
&
Co
Concentration, Conductivity
Mg * Alk (m«q/L> A Cond/100
Figure 8. Concentration profiles of calolum, magnesium, alkalinity
and profile of conductIvlties In Havana lowlands aquifer, 9/19/85.
D Ca2*. + Mg2*. O Alkalinity, ASpeolfio Conductance.
-------
o
v^
o
o
100
35
2OO
TIm« (days)
•f SO
300
OS
Figure 9. TiM aeries graph of oalolua concentrations
In Havana lowlands aquifer.
D 35. * 50. O 65 feet.
-------
I
25
24 -
23 -
22 -
21 -
2O -
10 -
18 -
17 -
•
10 -
15-
14 -
13 -
12 -
11 -
10
100
200
Tim* (day*)
300
4OO
Figure 10. Tiae series graph of magnesium concentrations
In Havana lowlands aquifer.
D 35. * 50. O65 feet.
-------
?
&
&
c
1=
o
JX
100
200
Ttm« (doy»)
300
Figure 11. Time aeries graph of alkalinities
in Harana lowlands aquifer.
D 35, * 50. O 65 feet.
-------
I
o
o
-10
-20
-30 -
-4O -
-50 -
-60 -
-70 -
-80 -
-90 -
-100 -
-110
0.0
Sulfote
0.1
0.2
i
0.3
Concentration (mM)
Chloride o Na
0.4
6. K * 10
0.5
Figure 12. Concentration profiles of sulfate, chloride, sodium,
and potassium in Havana lowlands aquifer, 9/19/85.
D SOj,2", + Cl", O Na*, A K*.
-------
100
I -
200
300
(days)
Figure 13. Time aeries;graph of sodium concentrations
In Havana lowlands aquifer.
D35, * 50, O65 feet.
-------
o
\
£
20 -:
15
100
300
Tlm« (days)
Figure 14. Time series graph of aulfate concentrations
in Havana lowlands aquifer.
D 35, + 50, O 65 feet.
48
-------
o>
o.o
0.8 -
O.7 -
0.6 •;
0.5 -
0.4 -
O.3
0 •
100 . 200
Tim* (day*)
300
Figure 15. Time aeries graph of potassium concentrations
In Havana lowlands aquifer.
D35, * 50, O 65 feet.
-------
•o
*c
3.
c.
u
100
200
Tim* (days)
Figure 16. Time aeries graph of chloride oonoentrationa
in Havana lowlands aquifer.
D35, + 50,O 65 feet.
'50
-------
ex
9
o
0
-10 -i
-2O -
-30 -
-4O -
-50 -
-60 -
-7O -
-80 -
-90 -
-100 -
-11O
7.5
i
7.7
7.9
8.1
PH
Figure 17. Profile of pH raluea in Havana lowlands aquifer, 9/19/85.
51
-------
a
a
-20 -
-JO -
-4O -
-50 -
-60 -
-70 -
-80 -
-9O -
-10O -
-110
O.OO
O.02
NftraU
0.04
O.OO
O.Od
O.10
Concentration (mi
•f • Phoapnat* * 100
Figure 18. Concentration profiles of nitrate and orthophosphate
oonoentratlona In Havana lowlands aquifer* 9/19/85*
D 1103-, * o-PO*3-.
: 52
-------
V—»
s
*->
z
10O
2OO
Tlm« (day*)
300
Figure 19. Time series graph of nitrate concentrations
In Havana lowlands aquifer.
D 35, * 50,O 65 feet.
53
-------
a.
M
o
.c
OL
O
200
Tim* (days)
Figure 20. Ti«e series graph of orthophosphate concentrations
in Havana lowlands aquifer.
D35, * 50, O65 feet.
-------
w«
1
v«*
I
c
o
o»
o
»•
o
Figure 21. TiM aeries graph of organic carbon concentrations
In Harana lowlands aquifer.
O 35, * 50, O 65 feet.
55
-------
600
Eh O2/H2O2
500
400
(35-65)
NOJTNHj
(35-65)
300
ui
200
100
— Fe(OH)3/Fe2+
(35-65)
Eh
— (104)
Fe(OH)3/Ft2*
— (104)
Figure 22. Comparison of" Measured and calculated redox potentials
in Havana Lowlands ground waters.
MuBbers in parentheses are depths. Dotted lines in Eh range
include extreaw values, solid line includes all but extreme values.
Arrows indicate lower bounds of calculated potentials.
56
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