Experiments Illustrating Metal Hydrolysis and Redox
Equilibria in Acid-Mine Waters
D. B. Levy* and W. H. Casey
oxidation-reduction equilibria, and mineral stability relaABSTRACT
tionships. The purpose of the field exercise was to integrate these concepts together. During the trip students
Teacherscan improveenvironmentalchemistrycourses by
record chemical measurements on a data sheet providapplyingthe conceptstaughtin the classroomto actual field
ed, and then answer a series of questions pertaining to
situations. Wedesignedan exerciseto emphasize
the interrelathe data. Analysis of the geochemicaldata in a final comtionships betweenacid-basechemistry,oxidation-reduction
posite exercise, in conjunction with visual inspection of
equilibria, andmineralsolubility in a field setting. Students
the site, allows the students to assess the impact of acid
measured
chemicalandphysical properties(i.e., pH,Eh, Fe
generation by mine tailings on local stream quality.
concentrations,
andelectrical conductivity)of boththe acidic
minedrainagein a large surfaceimpoundment
(pond),andthe
Theobjective of this article is to provide a specific exwaterat twolocationsin an adjacentstream.In a final analyample of howan integrative approach in a field setting
sis, comparisonsweremadebetweenproperties of the pond
can be used to synthesize those geochemical concepts
water and the stream water, upstreamand downstream
from
learned in the classroom. The concepts presented here can
the pond,by plotting the data on mineralstability diagrams
be easily adapted to other environmentalsettings. For exfor the common
Fe minerals--jarosite [KFe3(SO4)2(OH)6(s)], ample, this approach can be used to demonstrate the facpyrite [FeS2(s)], amorphous
Fe(OH)3(s)--and
Pb minerals-tors controlling chemical gradients from depth in lakes,
galena[PbS(s)],cerussite[PbCO~(s)],
andanglesite[PbSO4(s)].
wetlands, and groundwaters, or to assess the impact of
Thedatacollectedsupported
ourvisual observations
that acidic
surface runoff or industrial effluent on stream quality.
minewaterprecipitates Fe(OH)~(s)
andAI(OH)~(s)
as
els throughoverlyingstreamterrace deposits andmixeswith
adjacentstreamwater.Thisarticle presentsa specific example
GENERAL SETTING
of howan integrativeapproach
canbe usedto synthesizeseveral
importantgeochemical
conceptscommonly
taughtin the classThe abandoned Spenceville copper mine in central
room.Theexercisecanbe easily adaptedto otherenvironments,
California was chosen as the study location (Fig. 1). The
suchas salt marshes,lakes, or fresh-waterwetlands.
Spenceville site is severely eroded and contains a large
surface impoundment(pond) that fills with acidic runoff
originating from adjacent mine tailings and waste rock.
EACHERS OF ENVIRONMENTAL CHEMISTRY COURSES
The water in the pond is very acidic (pH = 2.5) and conoften employ the concepts of mineral weathering
tains elevated levels of Cu, Fe, and SO42-. The site was
and stability relationships, oxidation-reduction equilexcavated for Cu in 1865, and the iron-sulfide minerals,
ibria, and acid-base chemistry to describe geochemical
mostly pyrite (FeS2), were combustedto produce sulfurprocesses that occur in natural systems. Typically, the stuic acid (H2SO4). The ore body, which is now hidden
dents complete routine problem-solving assignments that
beneath the pond, is hosted by Jurassic-age marine volinvolve thermodynamicand equilibrium calculations for
canic rocks.
hypothetical chemistries in soils, sediments, and natural
The rocks are metamorphosedto greenschist grades
waters. While the students’ knowledgeis greatly enhanced
and deformedso that they dip at a high angle with respect
through problemsolving in the classroom, an integrative
to horizontal. This ore deposit is knownas a "Cypressmethod that emphasizes the practical applications of
type" massive sulfide deposit by analogy with deposits
those geochemical concepts is often superior. Levy and
on the island of Cypress in the Mediterranean Sea. These
Graham(1993) utilized a field setting to illustrate the
deposits were all formed during sub-sea-floor volcanism
interrelationships between soil chemical, physical, and
that was associated with the formation of new oceanic
biological properties in a landscape-scale exercise for an
crust (Barnes, 1979).
introductory soil science course. This integrative exercise
Without formal geologic training, the students have
was well-received by the students and faculty, and exdifficulty understanding the setting of the ore body and
tended the traditional soil science field exercises that
the subsequent reactions. Wetherefore drew an analogy
often failed to synthesize those concepts taught in the
betweenthe present-day ore body and pictures of the subclassroom.
sea hydrothermal vents (black smokers) that have been
Wedesigned a 4-h field exercise for an introductory
filmed with deep-sea submersible vehicles. The important
course in aqueous geochemistry that was conducted at
point is that the Spenceville ore minerals, and the sulfide
an abandoned copper mine. The lectures preceding the
minerals in the hydrothermal vents on the modern sea
field exercise separately covered acid-base chemistry,
floor, both precipitated in a sulfide-rich, oxygen-poorenvironment. The minerals are presently unstable, and
Department of Land, Air, and Water Resources, Univ. of California,
generating acidity, because the ore body has been excaHoagland Hall, Davis, CA 95616. Received 23 May 1994.
vated and exposed to rain, which has a muchhigher 02
*Corresponding author.
content than the original seawater or the local groundwaters (Fig. 2).
Published in J. Nat. Resour. Life Sci. Educ. 24:27-32 (1995).
T
J. Nat. Resour. Life Sci. Educ., Vol. 24, no. 1, 1995 ¯ 27
N
Spenceville
pit
100 m
Fig. 1. Location
of the Spenceville
Minestudyarea.
Oxidation of the sulfide ore deposits, upon exposure
to O2 and H20, releases H2SO4, as shown by the oxidation of FeS2:
FeS 2 + H20 + 7/2
FeSO4 + H2SO4
0 2 --
[1]
The ferrous iron (Fe2÷) released from this initial reaction is then oxidized by 02, resulting in the production
of ferric iron (Fe3÷). A catalytic cycle is generated,
whereby the Fe~÷ produced oxidizes more FeS2, generating additional FeE÷ and acidity. In pure systems, the
oxidation of Fe2÷ to Fe3÷ is considered to be the ratedetermining step in the reaction schemedescribed above.
In natural systems, however, the rate of oxidation of
FeE÷-- Fe~÷ is greatly accelerated by bacteria of the genus Thiobacillus (Singer and Stumm, 1970; Nordstrom,
1982).
MATERIALS
AND METHODS
Each student is provided with a 9-page handout that
includes a brief history of past mining activities at the
site, a general discussion on the chemistry of acid mine
deformation
ca. 180millionyearsago
drainage generation, Eh-pH stability diagrams for the
chief Fe and Pb minerals (e.g., van Breeman, 1982;
Faure, 1991), and a sheet of graph paper for the construction of an acid-base titration curve (Harris, 1991). The
handout also provides the students with a table for recording chemical data, in addition to a complete list of
questions (Table 1) that is completed during the 4-h
period. The students work as a group and are provided
with portable devices for determining solution pH, electrical conductivity (Orion Research Inc., Cambridge,
MA)and oxidation-reduction potential (Cole-Parmer,
Chicago, IL). The pH meter was calibrated in the field
using standard buffer solutions of pH 4.0 and 7.0.
Calibration of the Eh meter was achieved by immersing
the Pt electrode in Zobell’s solution prepared in 1.0 M
KCI, and adjusting the reading to + 192 mV(APHA,
1989). The students are also supplied with sampling bags,
a waterproof marker, clipboard, thermometer, water bottle, 50-mLdisposable beakers, a 1-mLpipette, 50 mLof
0.04 MNaOHsolution, and test strips for Fe2+ and total Fe measurements(Quantofix Iron-100, Aldrich Chemical Co., Milwaukee, WI).
The students first measure pH, Eh, electrical conductivity, and dissolved Fe concentrations in the pondwater,
and then in Little Dry Creek, both upstream and downstream from the Spenceville pond. Afterward, the students are asked to interpret the data collected in a final
composite exercise, which allows them to use their
knowledge of geochemistry to understanding the natural processes occurring at the mine site.
ancientseafloor
Pond Location
1860-1920
~
presentday
Fig. 2. Sequence
of eventsleadingfrom
theJurassic
sea-floorhydrotherreal activity to the present-day
ore bodythat is exposedand
weathering.
28¯ J. Nat. Resour.
Life Sci. Educ.,Vol. 24, no. 1, 1995
At the pond location we identified surface runoff and
lateral groundwater movementas the main transporting
agents of acidity and metals to the.pond water. The students measure the total Fe and Fez+ concentrations, in
addition to the pH, Eh, and conductivity of a surface
sample of the pond water. Previous studies of the Spenceville site have shownthat the upper 1 m of the pond is
oxygen-rich and well-mixed, whereas oxygen is virtually
absent below this depth, and the conductivity increases
with increasing depth in the pond. Theseincreases in conductivity becomeapparent by lowering the conductivity
cell to successively greater depths. The stratification of
the water column, and trends in the conductivity, temperature, and dissolved oxygen, were discussed in the context of steady-states of chemicalequilibrium as controlled
by mixing of the surface waters.
The methods used by early miners to remove Cu from
the orebodyillustrate the application of simple geochemical concepts. After the ore was excavated, piles of the
ore were roasted and then leached with water at the site.
Copper was removedfrom the leachate by precipitation
onto the surfaces of metallic Fe. The reaction is:
Fe°(s) + Cu2+(aq) -- Fe2+(aq) + Cu°(s)
To verify this reaction, the students lower an iron ring
into the pond to a depth of 5 mand inspect it later for
evidence of Cu precipitation onto the Fe surface. The
finely divided Cu metal, which is commonlyblack, appears as a coating on the metallic Fe surface.
The above exampledepicts a simple electron-exchange
reaction. Similarly, a separate class of hydrolysis reactions can be illustrated by raising the pH of samples of
the highly acidic pond water. The acidity in the pond
results from the production of hydroniumions (H30 +)
formed when the H2SO4produced in Eq. [1] reacts with
H20. The two dissociation equilibria are:
H2SO4(aq) + H20(1) -- H30+(aq)
+ HSO4-(aq)
(pK~ -~--3)
[3]
HSO4(aq) + H20(1) -- H~O+(aq) z- (aq)
(pK2 = 1.9)
[4]
Thus, a pond water pH equal to that of pK2 would yield
a solution containing equimolar concentrations of
HSO4-(aq) and SO42-(aq), according to the HendersonHasselbalch equation:
pH =
pK2
-)
q- log (SO42-/HSO4
[5]
At the observed pH (2.5) of the pond water, however,
the amountof HSO4-(aq)relative to SO42-(aq) in solution is small. Thus, the principle reaction that will govern
the changein pHof the pondwater as it is raised is given
by:
H30+(aq)
+ OH-(aq) -- 2H20(1
)
[6]
The pond water is roughly in equilibrium with ferric
hydroxide [Fe(OH)~(s)], a major weathering product
sulfide oxidation that can be easily recognized as a conspicuously orange material commonlypresent as rock
coatings. Uponadding a strong base to the pond water,
additional amounts of Fe(OH)~(s) form:
Fea+(aq) + 30H-(aq) -- Fe(OH)a(s)
[7]
Table 1. Representative student questions for the field exercise
in geochemistry.
Pond location
1. Write the twohalf-reactions that describe the precipitation of soluble
copper[as Cu2+laq~]onto the surface of native iron [Fe°ls~] to give an
overall electrochemical reaction.
2. Since the free energychange of an electrochemical reaction is° related
to the electrical work done on the system, we can write: AG =-nFE*.
Calculate z~G° for the net reaction given that E° = +0.78 V and F =
9.65 × 104 C mo1-1. Does the sign of ~G° agree with your observations? Explain.
3. Write the reaction for the titration of the pond water with the NaOH.
4. Howwouldthe titration curve shift if a weakeracid {i.e., higher pKa~
were being titratod?
5. Note the pHat which precipitates begin to form as the pond water is
titrated to higher pHvalues. Whatare the precipitates composedof?
Stream location
1. Did you note any differences regarding the appearance of the stream
upstream and downstreamfrom the Spenceville pond7
2. Howwouldthe availability of phosphorusto aquatic life in the stream
be affected by diurnal trends in photoreduction of FeIOH}3?
3. Howcan we be sure that the upstream location is representative of
"true" backgroundfor comparison with the downstreamlocation?
Minetailin~s location
1. Lookfor evidence of unoxidized pyrite and jarosite formation in the
waste material. Compareand contrast the luster, hardness, and color
of the two minerals and classify them as either primary or secondary
minerals.
2. List the chemicaland physical eharaetoristies of minespoil materials
that adversely affect plant growth and pose problemsin the revegetation of the site.
3. Write a balancedequation for the reaction of pyrito IFeS2Jwith oxygen
and water to form sulfuric add and soluble ferrous sulfate.
4. Given that oxygenand microorganismsare already present to someextent, whydo miningoperations cause an accelerated release of acids and
sulfate ~SO4~-Ifrom the ore?
5. Whattypes of amendmentswould y~u suggest adding to the tallinge
to aid in the establishmentof vegetative cover? What;other factor must
be considered whenrecommendingspoil amendmentsto control acidity
over long time periods?
Final composite exercise
1. Plot the Eh and pHof the pond water and the two streamwator locations on the stability diagramsprovided.
a. WhichPb and Fe minerals migh’~precipitate in the pond water?
b. WhichPb and Fe minerals might precipitato as water from the pond
movesinto Little DryCreek?Wasthere any visual evidencefor this?
Explain.
e. Whatreactions do you think controls the measuredEh of the waters?
Whatare the sources of error in this approachto redox equilibria?
2. Fromyour analysis of the data, what is your conclusion regarding contamination of Little Dry Creek from the Speneeville pond?
As the solution pHcontinues to increase, and essentially
all of the Fe in solution is precipitated as Fe(OH)3(s),
dissolved A13+(aq) begins to precipitate as the more
soluble AI(OH)3(s):
Ala+(aq)
+ 30H-(aq) -- AI(OH)3(s
)
[8]
To demonstratethe reactions that occur during the neutralization of acid minewaters, a series of beakers is prepared, and a titration curve is constructed by successive
additions of NaOHto equal volumes of acidic pond
water. The students observed an increase in the solution
pH, and precipitation
of Fe(OH)~(s), followed
AI(OH)3(s) as the titration progressed. The consumption
of OH-(aq) by Eq. [7] and [8] was discussed as
mechanismthat buffers the solution pH during the titration. After collecting the data, the students plotted the
titration curve on the graph paper provided (Fig. 3). Table 1 includes typical questions for this portion of the exercise.
J. Nat. Resour. Life Sci. Educ., Vol. 24, no. 1, 1995 * 29
a system is from equilibrium. For example, the solubility product for amorphous Fe(OHh(s) (Yariv and Cross,
1979) is given by:
10,
(Fe3+)(OH-)3 = Ksp = 2.5 x
2’
0
,
I
,
,
.
,
,
2
4
6
8
1
0
Volume of base added (mL)
Fig. 3. Titration of the acidic pond water (pH = 2.5) with 0.04 M
NaOH. The portions of the curve corresponding to neutralization
of H 3 0 and precipitation of amorphous Fe(OH)3(s) (4.[6] and
[71), and precipitation of AI(OH)3(s) (Eq. [8]) are shown.
+
Stream Location (Little Dry Creek)
The potential uses of the Little Dry Creek stream water
for domestic supply, irrigation, and migration habitats
were discussed at the stream location. The students visually assessed the potential for stream contamination by subsurface discharge of the pond water, and then measured
the pH, Eh, electrical conductivity, and Fe concentrations, both upstream and downstream from the pond
(Fig. 1).
After the general chemistry of the stream site was contrasted with the pond, several locations were identified
in the streambank where pondwater was seeping into the
higher-pH stream. These seepage sites are characterized
by blisters of amorphous precipitate consisting of
Fe(OH)3(s) (orange in color) surrounded by AI(OH)3(s)
(white in color) (Fig. 4). The students were asked to discuss the existence of these precipitates in terms of the
general hydrolysis chemistries of Fe3+ and Ai3+, resulting in the separation of Fe and Al in the streamwater (Fig.
5). The contrasting colors of the two precipitates were
then used to discuss the relative solubility of the minerals, and the concepts of supersaturation and undersaturation were introduced as a means of describing how far
[9]
Slow precipitation kinetics, however, may result in an ion
activity product (IAP) value [(Fe3+)(OH-)3] that is
greater than the Ksp(Le., IAP/Ksp > l), and the solution is said to be supersaturated with respect to amorphous Fe(OH)&). Here, a condition of supersaturation
means that additional precipitation of amorphous
Fe(OH)3(s) is thermodynamically possible. Alternatively, the solution may be undersaturated with respect to
amorphous Fe(OHb(s) if IAP/K,, < 1, and Fe(OH)3(s)
can potentially dissolve.
In a strong sense, the observed separation of A1 from
Fe via progressive neutralization of the pond water observed during this part of the exercise is directly analogous to the process of enrichment that formed the original
ore body. In this case, Cu, Fe, and Zn minerals separated from a homogeneous fluid to reach ore-grade concentrations in the rock. Typical questions for this portion
of the exercise are given in Table 1.
Mine Tailings Locations
The top of the tailings piles provides a complete overview of the setting and is an excellent location to discuss
erosion of mine tailings and deposition downslope. Here
we discuss the generation of acid from the mine tailings,
the problems associated with establishment of plant
growth on the tailings, and potential methods for the
remediation of acidic mine tailings. The students note the
colors of the various types of mine tailings present at the
site and then look for evidence of unoxidized pyrite
(FeS2) and secondary minerals, such as jarosite
[KFe3(S04),(OH),] and kaolinite [A12Si205(0H)4].We
emphasized the importance of jarosite in controlling the
remediation strategies. This mineral results from the
weathering of pyrite and buffers the soil acidity by releasing ferric iron as it dissolves:
Acid
3+
AI,
3+
Fe: pH
=
2
Less Acid
Near Neutral
Fig. 4. Photograph showing the precipitation of Fe and AI in Little
Dry Creek.
30
J. Nat. Resour. Life Sci. Educ., Vol. 24, no. 1 , 7995
Fig. 5. Steps leading to the separation of Fe3+ and AI3+ as amorphous stream precipitates of Fe(OH)3(s) and AI(OH)3(s).
Table2. Studentevaluationof the field exercise.
Rating
(Disagree} 1 2345 {Agree)
1. Thisfieldexercise:
a. Increasedmyunderstanding
of acid-basechemistry.
b. Increased
myunderstanding
of mineralstability relationships.
c. Broadened
myunderstanding
of oxidation/reduction
relationships.
d. Wassuccessfulin demonstrating
the interrelationships
between
acid-basechemistry,
redoxequilibria,and
mineralstabilityin naturalsystems.
e. Increasedmyknowledge
aboutthe formation
andremediation
of acid minespoils.
2. I wouldrecommend
this integrativeapproach
to otherinstructorsin their teachingof geochemical
concepts.
Number
outof eight responding.
KFe3(SO4)2(OH)6(s)+ (aq) -- K + ( aq
0t0044
00053
01124
00125
00035
00017
dents to design aa approach to revegetation of the mine
tailiings. This task leads into a discussion on the advantages and disadvantages of various types of soil amendments and plant species for the stabilization
of the
tailings, and methods for long-term monitoring of the
site. Typical questions for this portion of the exercise are
given in Table I.
+ 3Fe3+(aq) + 2SO42-(aq) + 3H20(1)
followed by:
Fea+(aq) ÷ 3H20(1) -- Fe(OH)3(s) ÷ 3H+(aq)
yielding the conspicuous orange Fe(OH)a(s) and causing
a net decrease in pH.
The students are asked to makea saturated paste with
the fine-textured spoil material and record the measured
pH on the data sheet. Wediscuss the chemical and physical properties of the mine spoil materials that present
difficulties for plant establishment, and then ask the stu-
Final Composite Exercise
The final compositeexercise allows the students to synthesize the geochemical information collected during
the field exercise. Here the students begin to understand
I
1"20 1 I
1.00
I
--,,,,,0~
""’-\,.
.80
Plattnerite
.6O
.4O
I Cerussite
I (PbCOa)
" --- Galena+ Sulfur I -
Eh
0
-.20
-.40
-.60 t
2
4
6
pH
8
10
12
z+,
Fig. 6. TheEh-pH
diagram
provided
in the studenthandout
for Fe
amorphous
Fe(OH)a(s)
jarosite,
and
pyrite
at
25
°C.
The
activities
, are: (Fe2+) + (Fe3+) = 10-4, SO42-= 5 ×
of solutionspecies
10 -3 M,andK= 5 x 10-4 M.Shaded
areasrepresent
solid phases.
Boundaries
areshown
for samples
fromthe pond(P) andthe Little
DryCreekstream(S) (diagram
after vanBreeman,
1982).
-.80
2
I
8
10
12
Fig. 7. TheEh-pH
diagram
provided
in the studenthandout
for lead
mineralsat 25°C.TotaldissolvedS = 10-1 M,(Pb) = 10 -6 M,
andPCO
are shownfor samplesfrom
2 = 10-s MPa.Boundaries
thepond
(P)andtheLittleDryCreek
stream
(S) (diagram
afterFaure,
1991).
J. Nat.Resour.
Life Sci. Educo,
Vol. 24, no.1, 1995¯ 31
the processes controlling mineral stabilities at the site,
and are in a position to assess the impact of acid generation from the mine tailings on the water quality of Little Dry Creek. First, the students are reminded of the
Fe(OH)3(s) precipitation observed as the pond water was
titrated to higher pH values. Then, students remember
observing orange precipitates in Little Dry Creek downstream from the Spenceville pond (Fig. 4). We discuss
the movement of the pond water into Little Dry Creek
as an analogous situation, whereby the acidic water is in
effect being titrated to higher pH values as it mixes with
stream water of higher pH. This final discussion is important to the overall success of the exercise because the
students are asked to link the results of experiments on
the pondwater with processes that also occur naturally.
The last step in the exercise is a compilation of measurements in the three environments. The students plot the
Eh and pH data from the pond and stream waters on an
Eh-pH diagram showing the stability fields for amorphous Fe(OH)3(s), jarosite, and pyrite (Fig. 6), and the
stability fields for anglesite [PbSO4(s)], cerussite
[PbCO3(s)J, and galena [PbS(s)] (Fig. 7). The data for
the pond water lie in the jarosite stability field shown in
Fig. 6, while the data points for the Little Dry Creek
waters lie in the stability field for Fe(OH)3(s). This indicates that precipitation of Fe(OH)3(s) becomes thermodynamically favored as the pond water mixes with the
stream water, a result that is consistent with both the experimental results and the observations along the stream
bank (Fig. 4). Similarly, Fig. 7 indicates that PbCO3(s)
is a solid phase that may control Pb solubility in Little
Dry Creek. Furthermore, the data collected indicate a
decrease in the stream water pH, and an increase in both
the total dissolved solids and Fe2+ concentrations of the
stream water downstream from the Spenceville pond.
These combined observations support the hypothesis that
the pond water is indeed moving into Little Dry Creek,
acidifying the water and loading the stream with dissolved
metals (e.g., Fig. 4). Representative questions for the final composite exercise are given in Table 1.
CONCLUSIONS
Field exercises such as these allow the students to synthesize geochemical concepts taught in the classroom and
to apply them to practical situations. The field experiments were found to provide a critical link between the
32
• J. Nat. Resour. Life Sci. Educ., Vol. 24, no. 1, 1995
classroom discussion and the natural processes of acid
generation and neutralization at the mine site. The students induce the very changes in chemistry that are observable to a trained eye in the field.
The response to our student questionnaire which was
designed to determine the effectiveness of the field exercise indicates that it was successful in demonstrating the
interrelationships between acid-base chemistry,
oxidation-reduction equilibria, and mineral stability
(Table 2). This integrative field exercise can be applied
to the teaching of geochemical concepts at other educational institutions as well. While a disturbed mine site may
not be close at hand, settings such as lakes or wetlands
are ideal environments for the application of this
technique.
ACKNOWLEDGMENTS
This research was funded in part by grant no. DEFG03-92ER14307 from the Office of Basic Energy Science at the U.S. Department of Energy. Special thanks
are extended to the 1994 students of Hydrological Sciences 134, Aqueous Geochemistry, at the University of
California in Davis for their enthusiasm and participation. We also thank Mr. Kit Custis for his assistance in
the field.