Aqueous-Sulfate Stable Isotopes—A Study 

of Mining-Affected and Undisturbed Acidic 

Drainage

By D. Kirk Nordstrom, Winfield G. Wright, M. Alisa Mast, Dana J. Bove, 

and Robert O. Rye

Chapter E8 of

Integrated Investigations of Environmental Effects of Historical 

Mining in the Animas River Watershed, San Juan County, Colorado

Edited by Stanley E. Church, Paul von Guerard, and Susan E. Finger

Professional Paper 1651

U.S. Department of the Interior

U.S. Geological Survey
Contents

Abstract .......................................................................................................................................................391

Introduction.................................................................................................................................................391

Acknowledgments ............................................................................................................................392

Purpose and Scope ..........................................................................................................................392

Previous Studies ........................................................................................................................................392

Inorganic Pyrite Oxidation ...............................................................................................................393

Processes Contributing to Sulfate Stable Isotope Compositions in Mineralized Areas ..........394

Generalized Isotope Mass Balances on Sulfate .........................................................................394

Biotic and Abiotic Oxygen Isotope Studies on Sulfate Formation from Reduced Sulfur 

Oxidation ...............................................................................................................................395

Field Area—Bedrock Mineralization and Alteration ...........................................................................397

Methods.......................................................................................................................................................399

Water-Quality Sampling Methods ..................................................................................................399

Analytical Methods for Dissolved Constituents ..........................................................................404

Analytical Methods for Oxygen and Sulfur Isotopes of Dissolved Sulfate .............................404

Analytical Methods for Oxygen and Sulfur Isotopes of Minerals ............................................404

Results and Discussion .............................................................................................................................404

Distribution of Stable Sulfate Isotopes for Mined and Unmined Areas ..................................405

Influence of Gypsum and Anhydrite on Water Chemistry ..........................................................407

Three Dominant Processes Controlling Water Quality ...............................................................408

Relationship Between δ18OSO4

 from Pyrite Oxidation and δ18OH2

O ..............................................410

Conclusions.................................................................................................................................................413

References Cited........................................................................................................................................413

Figures

 1. Map of Animas River watershed study area, showing sampling site locations 

and distribution of δ34S values of dissolved sulfate in surface water in and near 

five mineralized areas ..............................................................................................................398

 2. Histogram of δ34S values for pyrite and gypsum/anhydrite from Animas River 

watershed study area ..............................................................................................................406

 3–9. Graphs showing:

 3. δ34S of sulfate plotted against δ18O of sulfate for surface waters and for 

 gypsum and anhydrite mineral samples in study area ................................................407

 4. δ34S of sulfate plotted against δ18O of sulfate for surface waters, designated 

 by area from figure 1 .........................................................................................................408

 5. Dissolved calcium versus dissolved sulfate .................................................................409

 6. Molar ratio of calcium to sulfate plotted as a function of pH ....................................409

 7. δ34S of dissolved sulfate in streams, springs, and adits plotted against calcium:

 sulfate molar ratio ..............................................................................................................410

 8. δ18O of sulfate plotted against δ18O of H2O .....................................................................411

 9. δ18O versus δD of study area surface waters and precipitation collected 

 in Rocky Mountain region ................................................................................................412
Tables

 1. Chemical and isotopic data for non-mining affected and mining-affected samples 

collected in the Animas River watershed study area ..........................................................400

 2. Sulfur and oxygen isotopic data for sulfide and sulfate minerals in the study area ........405

Abstract

The combined results from major-ion chemistry and 

sulfur isotopes show three dominant processes affecting water 

quality: gypsum and anhydrite dissolution, pyrite oxidation, 

and calcite dissolution. This conclusion is based on more 

than one hundred samples including water, pyrite, gypsum, 

and anhydrite, collected in the Animas River watershed study 

area and analyzed for stable sulfur and oxygen isotopes (δ34S

in pyrite and δ34S and δ18O in aqueous sulfate and sulfate 

minerals).

Gypsum and anhydrite are shown to be dominantly 

hypogene (hydrothermal) in origin with heavy sulfur and 

oxygen isotopic compositions (δ34S SO4

=15 to 18 per mil 

and δ18O SO4

= –3 to 5 per mil). Pyrite is significantly lighter 

(δ34S= –7 to 2.5 per mil). Gypsum and anhydrite dissolution 

and pyrite oxidation are the dominant sources of dissolved 

sulfate in these waters. Distinct trends of mixing between 

water dominated by gypsum/anhydrite dissolution and 

water dominated by pyrite oxidation are inferred from the 

data.

A tendency for aqueous sulfate in water samples from 

unmined areas to have slightly higher δ18O relative to aqueous sulfate in water from mined areas is also apparent in the 

data, although there is considerable dispersion in the values, 

with several overlapping data points. Part of this enrichment 

is caused by differences in isotopic composition originating 

from aqueous sulfate derived from different locations and 

alteration types. Another possible reason for this small enrichment in δ18O in aqueous sulfate in pyrite oxidation-dominated 

waters might be the greater production of ferric iron in mined 

environments that leads to greater utilization of oxygen from 

the water molecule rather than molecular oxygen from air 

in the formation of aqueous sulfate. A distinct relationship 

between the δ18O SO4

 values for pyrite oxidation-dominated 

waters and the δ18O in water is not apparent, although the 

δ18O SO4

 values are generally in a range that agrees with laboratory experiments. Evaporation, mixing of pyrite oxidationdominated water with gypsum/anhydrite-dominated water, and 

dilution or mixing of water subsequent to pyrite oxidation are 

all important processes that can confound the interpretation of 

oxygen isotopic data from aqueous sulfate in mineralized and 

mined areas.

Introduction

The oxidation of pyrite is a complex hydrobiogeochemical process that adversely affects the water quality of receiving 

streams by producing acid waters with high metal concentrations. Sulfide ore extraction and processing, for gold, silver, 

copper, lead, and zinc production, enhances the rate of pyrite 

oxidation, increases the rate of acid production, and can cause 

severe environmental damage (Plumlee and Logsdon, 1999; 

Jambor and others, 2003). A considerable amount of research 

in the last decade has been applied to understanding, quantifying, and modeling the process of sulfide-mineral oxidation 

during weathering (Nordstrom, 2000; Nordstrom, 2004). 

The interpretation of stable isotope data of aqueous sulfate, 

SO2–

4 , derived from pyrite oxidation has been a challenging 

research topic and not well understood (van Stempvoort and 

Krouse, 1994; Taylor and Wheeler, 1994). Analyses of the 

stable isotope ratios of both sulfur and oxygen in sulfate have 

the potential to constrain the possible sources of sulfate and 

mechanisms of sulfide oxidation when stable sulfate isotopes 

are used in conjunction with the stable isotopes of water. For 

example, if pyrite is oxidized primarily by ferric iron, the 

oxygen in aqueous sulfate is derived from the water molecule 

as shown in equation 1:

 FeS2

 + 14Fe3+

2

2+ + 2SO2–

4 + 16H+

 (1)

Alternatively, if pyrite is oxidized by oxygen, only 1

/

8

 of the 

oxygen in aqueous sulfate is derived from the water molecule 

and the rest is derived from atmospheric oxygen as shown by 

equation 2:

 FeS2

 + 7/

2

O2 + H2O Fe2+ + 2SO2–

4 + 2H+

 (2)

Chapter E8

Aqueous-Sulfate Stable Isotopes—A Study of MiningAffected and Undisturbed Acidic Drainage

By D. Kirk Nordstrom, Winfield G. Wright, M. Alisa Mast, Dana J. Bove, and Robert O. Rye

 + 8H O 15Fe
The large difference between δ18O of air (+23.5ο

/

οο) and 

that of water (variable but generally –50 to +10 o

/

oo depending on distance from ocean, and on elevation, temperature, 

climate, latitude, and storm tracks) might provide a tool for 

distinguishing these two mechanisms. The very slow rate of 

oxygen isotopic exchange between sulfate and water (Lloyd, 

1968) except at exceptionally low pH and (or) higher temperature (Hoering and Kennedy, 1957; Seal and others, 2000; 

Seal, 2003) would preserve the isotopic signature. The idea 

that oxygen isotopes could distinguish between microbial 

and inorganic oxidation was proposed by Taylor and others 

(1984a); this idea has encountered criticism (Toran and Harris, 

1989), however, because the actual pathways of electron transfer are not well understood.

Equations 1 and 2 are also used to represent the indirect 

(equation 1) and direct (equation 2) mechanisms by which 

chemolithoautotrophic bacteria are thought to catalyze pyrite 

oxidation (Silverman, 1967; Sand and others, 2001). The 

indirect mechanism is one in which Fe3+ directly oxidizes 

pyrite and the role of the bacteria is to reoxidize the Fe2+ to 

Fe3+ in solution. In the direct mechanism, the bacteria adhere 

to the pyrite surface and directly oxidize the Fe2+ and the S–

 in 

pyrite by some unknown enzymatic mechanism. Although the 

direct bacterial mechanism does not require that 87 percent 

of the oxygen in the sulfate ion must come from air oxygen 

(equation 2), this mechanism does allow it. In the bacterial 

mechanism, air oxygen could either be reduced to the oxygen in sulfate or to the oxygen in water. Part of the difficulty 

in identifying where the oxygen in sulfate originates is in 

knowing the mechanism by which bacteria utilize oxygen as 

the electron acceptor. The bacterial mechanism is important 

because, abiotically, the oxidation rate by Fe3+ is significantly 

faster than by O2; without microbial catalysis it is difficult to 

see how molecular oxygen would contribute to the resultant 

sulfate ion. By contrast, stable isotopes of sulfur and oxygen 

have the potential to constrain sources. For example, many 

mineralized rocks also contain gypsum and anhydrite in addition to pyrite. The stable isotopes of sulfate might allow the 

discrimination between pyrite oxidation and gypsum/anhydrite 

dissolution as sources of the dissolved sulfate in ground water 

(for example, van Everdingen and Krouse, 1988; Nriagu and 

others, 1991). Stable isotope data on sulfate minerals whose 

parent aqueous sulfate formed from sulfide oxidation in natural, acidic weathering environments have been presented by 

Rye and others (1992) and Rye and Alpers (1997) in addition 

to the authors already mentioned.

Acknowledgments

Fieldwork and isotope determinations were made possible from the AML Initiative funding. The senior author is 

grateful to the National Research Program, WRD, for their 

support of this research. Thoughtful and thorough reviews by 

Robert R. Seal, II, Charles N. Alpers, and Craig Johnson are 

greatly appreciated and led to several improvements.

Purpose and Scope

This report presents and interprets sulfate stable isotope 

data for more than 100 water and mineral samples collected 

from the Animas River watershed study area in the San Juan 

Mountains near Silverton, Colo. The samples represent the 

consequences of weathering reactions in a variety of geologic 

and hydrologic conditions and cover a wide range of water 

composition (Mast and others, this volume, Chapter E7). The 

San Juan Mountains have undergone extensive hydrothermal 

alteration and mineralization (Bove and others, this volume, 

Chapter E3), allowing us the opportunity to collect samples 

from a variety of mineralized zones including both mined and 

unmined environments across a large area.

Study objectives were:

• To review the literature on pyrite oxidation and the use 

of sulfate stable isotopes for hypotheses and arguments 

from theory and experiments that would assist in the 

interpretation of field data

• To collect a large enough data set so that the factors 

controlling the isotopic signature of dissolved sulfate 

in these waters could be discerned, especially the 

origin of the sulfate oxygen

• To determine whether a different isotopic signature 

exists for acid drainage from mines than acid drainage 

from rocks undisturbed by mining

• To improve our understanding of pyrite oxidation 

through the application of isotopes.

Previous Studies

The complex pathways by which the sulfur in pyrite 

becomes sulfate in aqueous solution involve several intermediate steps that may be affected by temperature, pH, solution 

composition, and microbial metabolism and ecology (Lowson, 

1982; Nordstrom and Southam, 1997; Nordstrom and Alpers, 

1999; Nordstrom, 2000). The oxygen isotope composition of 

the sulfate is affected by the isotopic composition of the reacting substances (Krouse and others, 1991; van Stempvoort and 

Krouse, 1994). The sulfur moiety changes from an operationally defined oxidation state of 1– to 6+, or a transfer of seven 

electrons per sulfur atom. This electronic transfer cannot happen in a single step (Pearson, 1976) and, consequently, several 

intermediate sulfoxyanions are implicated during the oxidation 

(Nordstrom, 1982; Goldhaber, 1983; Moses and others, 1987). 

Several studies have brought insight to these pathways from 

chemical laboratory experiments, isotopic laboratory experiments, quantum-mechanical calculations, and field data. This 

section reviews these pathways and formulates hypotheses 

about the source of oxygen in the final sulfate ion that is 

needed for the interpretation of isotopic data of dissolved sulfate in samples from the Animas River watershed study area. 

392 Environmental Effects of Historical Mining, Animas River Watershed, Colorado
Aqueous-Sulfate Stable Isotopes and Acidic Drainage 393

This review emphasizes recent investigations that reveal the 

nature of bonding of oxygen, whether from molecular oxygen 

or from water oxygen, to the sulfur in the pyrite surface and 

to reaction products such as thiosulfate and sulfite.

Inorganic Pyrite Oxidation

The initial attack on pyrite under oxic conditions and in 

the presence of water is a surface reaction that promotes a dissociative sorption of the water molecule at the pyrite surface, 

forming iron-oxygen bonds across the surface that are precursors to hydrous ferric oxide formation (Rosso and others, 

1999). Iron oxyhydroxide surface formation clearly precedes 

sulfate formation (Nesbitt and Muir, 1994). The form of iron 

oxyhydroxide has been identified as goethite (at pH 9.2), and 

the coordination and bond distances have been measured at 

the pyrite surface after small amounts of reaction (England 

and others, 1999). Pseudomorphs of goethite after pyrite are 

well known and seem to have formed slowly at circumneutral 

to high pH. When surface coatings of hydrated ferric oxide 

are formed, ferrous iron in this oxide layer can more readily 

reduce molecular oxygen than the ferrous iron in pyrite can 

(Eggelston and others, 1996). Hence, experimental evidence 

indicates that molecular oxygen reacts initially with the iron 

and not with the sulfur in pyrite.

One of the consequences of iron-oxygen bond formation is to depopulate sulfur surface sites so that they become 

slightly more hydrophilic and more susceptible to the formation of sulfur-oxygen bonds (Rosso and others, 1999). The 

origin of the oxygen combining with the sulfur appears to be 

from water unless the primary hydration sphere of the oxidized iron transfers its oxygen to sulfur when Fe3+ oxidizes 

a reduced sulfur species (for example, Moses and Herman, 

1991). The mechanism proposed by Moses and others (1987) 

involving the formation of a hydroxyl free radical from the 

hydration sphere of Fe(III) (the oxidant in experiments at low 

pH) has been experimentally verified by Bourda and others 

(2001, 2003) and further confirms that the oxygen that initially 

bonds to the iron and the sulfur is derived from the oxygen in 

water, not the molecular oxygen from air.

The oxidation of the sulfur, after the initial stage, is more 

complex because of the multiple sulfur species of intermediate 

oxidation state between 1– and 6+ that can form as proposed 

by Nordstrom (1982). Sato (1960) and Biegler and Swift 

(1979) used electrochemical measurements on pyrite in aqueous solution to show ferrous ions and the formation of either 

disulfide molecules or elemental sulfur. Bergholm (1955) 

was one of the first to carefully monitor pyrite oxidation in 

laboratory experiments and showed the formation of elemental 

sulfur and sulfate at temperature in excess of 50°C. One of 

the first studies to indicate thiosulfate formation on oxidizing pyrite is that of Steger and Desjardins (1978). Nesbitt and 

Muir (1994) monitored the initial reaction of pyrite with air 

and water vapor with X-ray photoelectron spectroscopy and 

found that the species disulfide, monosulfide, and polysulfide 

evolved through time (monosulfide decreasing, disulfide and 

polysulfide increasing) to eventually form sulfate and thiosulfate. Since then, Goldhaber (1983), Moses and others (1987), 

and Schippers and others (1996) have shown that thiosulfate, 

sulfite, and polythionates are all formed in different amounts 

depending on pH, oxidant, microbial ecology, and time.

Although the formation of intermediate sulfoxyanions 

has been measured during pyrite oxidation, the exact mechanism is still not entirely resolved. The consistent appearance of thiosulfate early in the reaction (those at low pH and 

without ferric ion as an oxidant) prompted Moses and others 

(1987) and Luther (1987, 1990) to postulate that a dioxygen 

molecule reacts directly with a disulfide molecule to form 

thiosulfate as the first dissolved sulfur oxidation product. At 

circumneutral pH, it was postulated that the dioxygen molecule reacted with reduced iron and then the oxidized iron provided an electron sink for electrons to leave reduced sulfur. At 

low pH, ferric ion can oxidize pyrite under anoxic conditions 

(Garrels and Thompson, 1960) so that all of the oxygen in 

sulfate must come from the water molecule.

An important question is whether air oxygen or water 

oxygen is used during the oxidation of sulfoxyanion intermediates to sulfate. Thiosulfate oxidizes to tetrathionate 

in the presence of pyrite (Xu and Schoonen, 1995) or Fe3+

(Williamson and Rimstidt, 1993), and tetrathionate tends to 

decompose through a series of reactions that include the production of trithionate, pentathionate, trisulfane monosulfonic 

acid or disulfane monosulfonic acid, and sulfite before ending 

with sulfate (Schippers and others, 1996; Druschel, Hamers, 

and Banfield, 2003; Druschel, Hamers, and others, 2003). 

Apparently, the only stages of the reaction that might require 

molecular oxygen are in the decomposition of tetrathionate 

to sulfate and in the oxidation of sulfite to sulfate. Druschel, 

Hamers, and Banfield (2003) suggested that the overall oxidation of tetrathionate in the presence of oxygen and ferric ion 

can be represented by:

 S4

O6

2– + 3Fe3+ + 11/

4

O2 + 9/

2

H2O 3Fe2+ + 4SO2–

4 + 9H+

 (3)

If it is assumed that all of the oxygen in tetrathionate is 

derived from the water molecule, then reaction 3 would dictate 

that about 34 percent of the oxygen in sulfate could be from 

molecular oxygen. This amount of molecular oxygen is in the 

range that Gould and others (1989) found in their isotopically 

controlled sulfide mineral oxidation experiments. Tetrathionate 

oxidizes through microbial mechanisms that generally do not 

require molecular oxygen, and the breakdown of polythionates by known enzymes is possible (de Jong and others, 1997; 

Friederich, 1998) in Acidithiobacillus ferrooxidans and Acidithiobacillus thiooxidans (formerly Thiobacillus ferrooxidans 

and Thiobacillus thiooxidans, Kelly and Wood, 2000). The 

most recent review summarizing information on the enzymatic 

oxidation of elemental sulfur to sulfate can be found in Kelly 

(1999). Polythionates are more stable than sulfite (Druschel, 
Hamers, and Banfield, 2003; Roy and Trudinger, 1970). 

Sulfite is another decomposition product of both thiosulfate 

and polythionates and the immediate precursor to sulfate. The 

enzymatic decomposition of tetrathionate produces thiosulfate and elemental sulfur (de Jong and others, 1997), both of 

which will further oxidize to sulfate. As mentioned previously, 

Schippers and others (1996, 1999) have outlined complex 

microbial pathways for the oxidation of tetrathionate, polythionates, polysulfane monosulfonic acids, and thiosulfate to sulfate. Recent work by Druschel, Hamers, and others (2003) has 

shown that the formation of hydroxyl radicals, known to occur 

on the pyrite surface in the presence of water, can oxidize 

polythionates much faster than molecular oxygen or Fe(III). 

They concluded that reaction with hydroxyl radicals is the 

only inorganic pathway for polythionates to oxidize to sulfate. 

The reaction appears to involve a trisulfane monosulfonic acid 

intermediate that can then further decompose in the presence 

of molecular oxygen to produce sulfur monoxides and sulfite 

and finally sulfate. These reactions would only allow molecular oxygen to be accommodated in the final sulfate by oxidation of sulfite to sulfate.

Sulfite oxidizes rapidly to sulfate in the presence of 

oxygen (Ermakov and others, 1997), but it oxidizes more 

rapidly by 2- or 3-fold in the presence of ferric ion (Brandt 

and van Eldik, 1998; Millero, 2001). Oxidation by ferric ion 

would not require any molecular oxygen, although studies by 

Holt and others (1981, 1983) indicate stoichiometric contributions of molecular oxygen during sulfite oxidation without 

iron. The amount of molecular oxygen may well be a function of pH, wherein more molecular oxygen appears in the 

sulfate at higher pH because of the insolubility of ferric iron 

(Seal, 2003). Furthermore, the oxygens in sulfite are known 

to exchange quickly with water oxygen (Hall and Alexander, 

1940; Betts and Voss, 1970) so that sulfite might only contain 

water oxygen regardless of its previous history. Hence, if the 

last stage of sulfur oxidation always involves the oxidation of 

sulfite to sulfate, then the primary pathway for molecular oxygen to enter the sulfate ion is during this reaction. The amount 

that enters will depend on pH, Fe(III) concentrations, oxygen 

concentrations, and possibly temperature.

Processes Contributing to Sulfate Stable Isotope 

Compositions in Mineralized Areas

Hydrothermal sulfide and sulfate minerals in mineral 

deposits generally have sulfur isotopic compositions that vary 

in the range of –10 to +25 per mil (Clark and Fritz, 1997; 

Faure, 1986; Krouse and Mayer, 2000; Seal and others, 2000) 

with sulfide minerals commonly in the range of –10 to +10 per 

mil and sulfate minerals having values significantly higher. 

Little fractionation of sulfur appears to take place during 

weathering to produce aqueous sulfate. Field (1966) hypothesized that the δ34S of sulfides would change little on oxidation to sulfate and reported data that suggested fractionation 

effects of –2 to +3 per mil. This conclusion has been supported 

by studies of supergene alunite (Rye and others, 1992) and 

jarosite (Rye and Alpers, 1997) from numerous ore deposits. 

Toran and Harris (1989) summarized most of the extant literature and confirmed, from both lab studies and field studies, 

that not more that a few per mil fractionation takes place during oxidation of S2

2– in pyrite to S6+ in dissolved sulfate.

The stable isotopes of sulfate should be a useful constraint for interpreting the operative processes of pyrite oxidation, especially at mining sites, if different source minerals 

had different sulfur isotope compositions and if the reaction 

pathways of molecular versus water oxygen are reflected 

in the oxygen isotopes of sulfate. Two complications make 

these interpretations challenging, however. First, there can 

be four general sources of sulfate in addition to pyrite oxidation: sulfate mineral dissolution, atmospheric sulfate, residual 

sulfate after partial reduction, and recycled sulfate (re-oxidized 

after reduction). Hypogene sulfate minerals are commonly 

associated with many sulfide deposits and need to be considered in any isotope study (Toran, 1987). Sulfate mineral 

dissolution, especially from gypsum and anhydrite, will be 

shown to be a major process in the Animas River watershed. 

In our study area, atmospheric sulfate in precipitation is very 

low in concentration relative to that found in most acid water 

and can be considered negligible. Cycling of sulfur through 

redox reactions is critical for some coal mine drainage situations and cannot be discounted in our study, but it seems 

unlikely to be a major process affecting sulfate stable isotopes 

in mountainous metal-sulfide deposits of the western United 

States. Water samples were collected in areas of high elevation 

and steep topography, and drainages did not infiltrate large 

wetlands. Small, localized wetlands were encountered in a 

few places and were not likely to have had significant effects 

on the sulfate stable isotopes. Second, an exchange of 18O–16O 

between aqueous sulfate and water can occur where special 

hydrologic conditions permit long residence times in low-pH 

environments (Chiba and Sakai, 1985; Alpers and others, 

1992; Seal and others, 2000; Seal, 2003). But the samples 

in this study are unlikely to have sufficiently low pH and 

sufficiently long residence time to permit isotope exchange. 

Evidence of long-term ponding is not evident in the water 

stable isotope data.

Generalized Isotope Mass Balances on Sulfate

Several possible mechanisms exist by which both atmospheric oxygen and water oxygen can be incorporated into the 

final sulfate from pyrite oxidation. Assuming the stoichiometry of reactions 1 and 2, and assuming contributions from the 

dissolution of sulfate minerals, the contributions of each type 

of oxygen will follow the isotope mass balance equation:

δ18OSO4

 = Y{X(δ18Ow + εw) + (1–X)[0.875(δ18Oa + εa) 

 + 0.125(δ18Ow + εw)]} + (1–Y) Σm xm(δ18Om) (4)

394 Environmental Effects of Historical Mining, Animas River Watershed, Colorado
Aqueous-Sulfate Stable Isotopes and Acidic Drainage 395

where X=the fraction of sulfate oxygen from the water (w) 

molecule by reaction 1 with the remainder (1–X) coming from 

air (a) oxygen from reaction 2, and with εw and εa

 the fractionation factors for water and air, respectively (van Everdingen 

and Krouse, 1985; van Everdingen and others, 1985; van 

Everdingen and Krouse, 1988; Toran and Harris, 1989; Clark 

and Fritz, 1997) and Y=the fraction of pyrite-derived sulfate 

with the remainder (1–Y) coming from the sum of the fractions (x) of dissolved sulfate minerals (m) and their isotopic 

compositions. As long as no oxygen isotopic exchange takes 

place between dissolved sulfate (from minerals) and the water 

molecule, there will be no fractionation factor. The complicating factors are

1. the mechanism by which microbes affect the transfer 

of oxygen

2. the mechanism by which unstable or metastable intermediate sulfoxyanion formation affects the isotopic 

composition

3. the variability in δ18OH2

O from climatic conditions, and

4. the extent to which δ18OH2

O actually represents the composition of the water that reacted with the pyrite at the site 

of oxidation (that is, complications from evaporation and 

from mixing of waters of different δ18OH2

O after oxidation 

occurred).

Equation 3 is linear and can be rearranged and simplified to 

give an empirical function of δ18OH2

O:

δ18OSO4

= m(δ18OH2

O) + b (5)

where the contributions from soluble sulfate minerals can 

be neglected, the atmospheric oxygen contribution is considered constant or can be determined by difference after 

the water-oxygen dependence has been ascertained, b is an 

empirical intercept, and where the fraction of water oxygens, m=0.875X+0.125. Hence, a plot of δ18OSO4

against 

δ18OH2

O could be solved for the proportion of reaction 1 versus 

reaction 2. Alternatively, if the fractionation factors are known, 

a diagram can be constructed with lines for varying proportions of these reactions and field data can be plotted on them. 

Both approaches have been utilized in previous studies (for 

example, van Everdingen and others, 1985; van Everdingen 

and Krouse, 1988).

Biotic and Abiotic Oxygen Isotope Studies 

on Sulfate Formation from Reduced Sulfur 

Oxidation

Several different lab experiments using O2 or H2O of 

known isotopic composition help to constrain the reaction 

mechanisms of sulfate formation. Lloyd (1967) found that sulfate formed from the bacterial oxidation of hydrogen sulfide 

incorporated 64–76 percent of its oxygen from water and the 

remainder from atmospheric oxygen. No significant fractionation was observed when water oxygen was incorporated, but 

a fractionation of –8.7 per mil was observed for atmospheric 

oxygen.

A strong dependency of sulfate δ18O on the δ18O of the 

water molecule was found in several experiments involving sulfur or sulfide mineral oxidation. Mizutani and Rafter 

(1969) oxidized elemental sulfur in two waters of different 

oxygen isotope composition using wet silty soil as a bacterial 

source. The results showed no fractionation, that is, δ18OSO4 had the same value as δ18OH2

O.

Microbial oxidation of reduced sulfur compounds has 

been investigated for many decades, and yet the biochemical processes by which sulfur-oxidizing autotrophic bacteria 

gain energy are not clearly elucidated. Depending on what the 

processes are, atmospheric oxygen may or may not be incorporated into the final sulfate ion. Microbial sulfur-oxidizing 

activity is caused by one or more enzymes, generally called 

sulfur (or sulfide) oxidases. An oxidase is any enzyme that 

catalyzes an oxidation reaction. If an enzyme catalyzes the 

incorporation of O2

 into organic compounds (which then may 

contribute the oxygen to the formation of sulfate), then it is 

known as an oxygenase. Much of the research on the biochemistry of sulfur- and sulfide-oxidizing microbes has been 

directed toward characterization of these enzymes and whether 

they behave as oxygenases (Kelly, 1982, 1999). The general 

picture of oxidation of sulfide and sulfur by acidithiobacilli 

(Madigan and others, 2000) is that the reduced sulfur combines with a sulfhydryl group of the cell, such as glutathione, 

which is then oxidized to sulfite (with or without molecular 

oxygen) by a sulfide oxidase. The sulfite can further oxidize 

to sulfate by a sulfite oxidase. Electrons are captured from the 

electron donor, such as pyrite or sulfur, and then transferred 

through an electron transport system (ETS) to the terminal 

electron acceptor, such as oxygen. The ETS is located in the 

membrane and includes a sequence of cytochrome enzymes 

which have been identified and studied in several thiobacilli. 

Adenosine triphosphate (ATP), the primary energy carrier 

in living organisms, is synthesized from adenosine diphosphate by the ETS. If the dominant sulfide-mineral-oxidizing 

bacteria operate with an ETS, then one would expect δ18OSO4 to be unaffected by δ18O2

 of air. With oxygenase, no metabolic 

energy is gained but δ18OSO4

 is affected by isotopic composition of atmospheric oxygen.

Suzuki (1965) investigated the oxidation of sulfur by 

Acidithiobacillus thiooxidans and found that the resulting 

sulfate incorporated no 18O from the air. A small amount 

of atmospheric oxygen was found when he used cell-free 

enzyme extracts, leading him to tentatively identify the sulfuroxidizing enzyme as an oxygenase. Numerous enzymatic studies, reviewed most recently by Kelly (1999), have attempted 

to better characterize sulfur oxidation pathways among the 

thiobacilli. Oxygenases have been clearly demonstrated to 

be active in Archaea such as Acidianus brierleyi, a hypothermophile and acidophile. An important point was made that, 

although thiobacilli have in common the oxidation of reduced 
sulfur to sulfate, different species appear to have different 

pathways. Kelly’s review did not settle the relative importance 

of oxygenase and the ETS, but it did give lines of evidence 

indicating alternative explanations. He also pointed out that 

the use of oxygenase during sulfur moiety oxidation would 

reduce the possible sources of metabolic energy gain but it 

would increase the apparent efficiency of the energy by 2- to 

3-fold. A revealing statement from Kelly (1999) provides a 

method for resolving the dilemma: “As yet, no attempt appears 

to have been made to determine the contribution in vivo of a 

sulfur oxygenase by providing bacteria with 18O-labeled oxygen or water to determine the relative contributions of each to 

the oxygen recovered in the sulfate produced by oxidation of 

sulfur (and thiosulfate). It is possible that exchange reactions 

might obscure the involvement of any direct oxygenation.” 

For the case of sulfide mineral oxidation, such studies have 

been reported but only rarely with the inclusion of controlled 

microbial catalysis.

Schwarcz and Cortecci (1974) found no correlation 

between the δ18O of the dissolved sulfate and that of spring 

and stream waters of northern Italy and concluded that evaporite mineral sources, pyrite oxidation occurring at different 

temperatures and pH values, and isotopic fractionation of the 

molecular oxygen were all confounding factors that could 

cause the δ18O SO4

 to vary independently of δ18O H2

O. They 

did, however, show that pyrite oxidation in water of varying δ18O caused a correlated increase of 0.6 per mil in the 

δ18O SO4 for every per mil increase in δ18O H2

O. This study was 

one of the first to demonstrate that at least half of the oxygen 

in the sulfate was derived from the water molecule. Bailey 

and Peters (1976) conducted pyrite oxidation experiments 

with 18O-labeled water at temperatures of 85°–110°C and 

found that 73–100 percent of the sulfate oxygen came from 

water oxygen. Taylor and others (1984a) found evidence for 

increases in δ18OSO4

 with increasing in δ18OH2

O in acid mine 

waters from the western United States, but possible contributions from other sulfate mineral sources were not considered. 

Between 23 and 100 percent of the sulfate oxygen apparently 

originated from water oxygen in the experiments of Taylor and 

others (1984b) depending on the conditions of the experiments 

(anaerobic versus aerobic, sterile versus inoculated with Acidithiobacillus ferrooxidans). Van Everdingen and others (1985) 

estimated that 37–74 percent of the sulfate oxygen from various field locations was derived from water. Van Everdingen 

and Krouse (1988) also found 29–100 percent of the sulfate 

oxygen derived from water oxygen in lab experiments and 

35–90 percent in field studies. They showed a plot of δ34SSO4 versus δ18OSO4

 that suggested a mixing line involving dissolution of anhydrite, and when combined with the known occurrence of anhydrite at their field site (Pine Point, Northwest 

Territories), the isotopic data strongly implicate anhydrite as 

contributing to the more positive δ18OSO4

 values. Cunningham 

and others (2005) showed similar mixing of aqueous sulfate 

derived from oxidation of pyrite and the dissolution of gypsum 

and alunite.

Gould and others (1989) demonstrated the consistent 

incorporation of water oxygen for a series of six experiments 

with four sulfide minerals (pyrite, chalcopyrite, sphalerite, and 

pentlandite) in nine different initial compositions of δ18OH2

O, 

pH=2.5, and with A. ferrooxidans. The percent water oxygen 

was 23–80 overall but averaged 66–71 percent for the pyrite 

runs. They noted that an increase occurred in the proportion of incorporated water oxygen with time, and that the 

molecular oxygen was likely incorporated into a sulfoxyanion 

intermediate during the reaction. Further work by the same 

group (Krouse and others, 1991) demonstrated clearly that O2

is incorporated into SO4

2– during sulfide mineral oxidation. 

They only looked at the oxidation of pentlandite (Fe,Ni)9S8, 

a monosulfide; these results, thus, may not be comparable 

to pyrite oxidation because it is recognized that monosulfides oxidize differently in acid solutions than disulfides do. 

Monosulfides readily form H2S, so that the oxidation process 

is similar to the oxidation of H2

S. Oxidation of H2S under 

aerobic conditions proceeds by reacting with molecular 

oxygen (Zhang and Millero, 1994). Furthermore, as pointed 

out by Krouse and others (1991), even though molecular 

oxygen is needed to oxidize sulfite (an intermediate during 

the oxidation of any reduced inorganic sulfur compound), 

there is rapid 18O–16O exchange between H2

O and SO3

2– (Betts 

and Voss, 1970; Eigen and others, 1961; Lloyd, 1967). The 

factors that control the isotopic exchange rate between these 

species are not well known, but Holt and others (1981) found 

that atmospheric oxygen contributed typically 25 percent or 

one of the four oxygen atoms in the sulfate molecule. Rapid 

isotopic exchange with water accounts for the remainder. 

These results could explain the fairly consistent trend of 

about 25–40 percent molecular oxygen occurring in sulfate 

from sulfide oxidation studies (Taylor and Wheeler, 1994; 

van Stempvoort and Krouse, 1994; Krouse and Mayer, 2000).

Reedy and others (1991) determined the proportion 

of actual isotopomers (S16On

18O4–n) produced during pyrite 

oxidation at pH 1 and 70°C and at pH 7 and 20°C by using 

vibrational spectroscopy and pure 18O2

 and pure H2

16O (and 

the reverse, pure 16O2

 and pure H2

18O) for initial reactant compositions. Runs with and without Fe3+ were completed. They 

found that the water oxygen dominated the composition of the 

sulfate but that some atmospheric oxygen was present; that 

is, when they used pure H2

16O, the sulfate was dominated by 

S16O4

 but some S16O3

18O was found. They also pointed to the 

oxidation of sulfite by O2

 and rapid exchange of sulfite oxygen 

with water oxygen. Consistent with other studies, they also 

found that the longer the experiments were run, the more water 

oxygen was found in the sulfate. Reaction kinetics clearly play 

an important role in the isotopic composition of oxygen in 

sulfate derived from pyrite oxidation.

The implication from these studies is that sulfite or 

polythionates are potentially the dominant, oxygen-isotopecontrolling intermediates in the overall oxidation reaction. In 

a review of atmospheric studies of oxygen isotope behavior 

396 Environmental Effects of Historical Mining, Animas River Watershed, Colorado
Aqueous-Sulfate Stable Isotopes and Acidic Drainage 397

during the air oxidation of SO2

 by Cunningham and others 

(1984), the isotopic composition of aerosol sulfate formed 

from the oxidation of SO2 is composed of one air oxygen and 

three water oxygens. The lab and field results of atmospheric 

sulfate formation demonstrate clearly that it makes no difference what the oxygen isotope composition of the SO2

 is, the 

composition of the oxidation product, SO2–

4 , is composed 

of one air oxygen and three water oxygens (Holt and others, 

1981). This conclusion favors the reactions,

 SO2(g) + 1⁄2O2(g) = SO3(g) (6)

and

 SO3(g) + H2

O(l) = H2SO4(aq) (7)

that results in the isotopic relation

δ18OSO4

 = 0.75(δ18OH2

O)+k (8)

where k is a semi-empirical constant that can be derived 

from equation 4 if all the factors are known. If the sulfite 

ion (hydrolyzed SO2

) is a necessary intermediate to SO2–

4

formation during pyrite oxidation, and one oxygen in sulfite 

is always derived from molecular oxygen, then a constant 

75 percent ratio of water oxygen might be expected. Note that 

in the sulfide mineral oxidation experiments of Gould and 

others (1989), the SO2–

4 approached 75 percent water oxygen 

at the longest times for four of the six runs. Although many 

of the experimental studies give results in this range, they do 

not consistently occur in this range. In the oxidation of pyrite, 

however, SO2(g) is not the main product formed, only aqueous sulfite ion in the presence of water, so the question is: 

by what pathway is sulfite formed? During H2S oxidation by 

oxygen, an SO2

 radical is formed which then hydrolyzes and 

oxidizes, forming thiosulfate, sulfite, and ultimately sulfate 

(Zhang and Millero, 1994). If thiosulfate forms directly from 

the attack of molecular oxygen on the disulfide in pyrite, then 

there should be two oxygen atoms derived from atmospheric 

oxygen. The second alternative is, if rapid exchange of water 

oxygen with the sulfite ion occurs, then there should be only 

one atmospheric oxygen atom. The third alternative is, if ferric 

ion dominates the oxidation of sulfite, that there should be no 

atmospheric oxygen atoms. These limits might provide the 

actual controls on the sulfate-oxygen isotopic composition. 

That is, the range of water oxygen in sulfate oxygen would be 

50–100 percent. This range is observed for experimental studies and for many field studies.

Sulfite ion oxidizes to sulfate rapidly in the presence of 

oxygen:

SO2–

3  aq + 1⁄2O2(g) = SO2–

4  aq (9)

Studies by Schippers and others (1996), Schippers and 

others (1999), Sand and others (2001), Druschel, Hamers, 

and Banfield (2003), and Druschel, Hamers, and others 

(2003) have shown that multiple complex pathways are 

likely involved with the intermediate stages of oxidation 

of the sulfur moiety from pyrite to sulfate. These pathways 

include formation of thiosulfate, polythionates, elemental 

sulfur, disulfane-monosulfonic acid, trisulfane-monosulfonic 

acid, and sulfite. These intermediate sulfoxyanions are likely 

to exchange their oxygen rapidly with water oxygen, but 

polysulfane-monosulfonic acids, like sulfite, might incorporate 

atmospheric oxygen during oxidation to sulfate. This possibility could also lead to larger δ18OSO4

 values.

A brief recapitulation is needed. Whereas lab studies 

indicate some consistency in the relationship between δ18OSO4

and δ18OH2

O, applying this information to field studies may 

be hindered by three important aspects. First, as pointed 

out by several researchers (Schwarcz and Cortecci, 1974; 

van Everdingen and others, 1985; Krouse and Mayer, 2000), 

dissolution of sulfate minerals, especially gypsum and anhydrite, must be quantitatively determined if they are part of the 

sulfate pool. Second, the H2

O in which the sulfate is dissolved 

may not have been the ambient H2

O at the time of pyrite 

oxidation. Evaporation and mixing with different waters can 

change δ18OH2

O without affecting δ18OSO4

. Third, oxygen isotope exchange can occur between SO2–

4 and water in low-pH 

waters at elevated temperatures (Chiba and Sakai, 1985). 

Without a detailed analysis of the hydrology and mineralogy 

of every site, the interpretation of the isotopic data may not be 

well enough constrained. The additional possibility of cyclic 

reduction and re-oxidation in areas of bogs, wetlands, soils, or 

collapsed anoxic mines might also confuse the interpretation.

Field Area—Bedrock Mineralization 

and Alteration

The Animas River watershed study area is underlain 

primarily by intermediate volcanic rocks of Tertiary age that 

are hydrothermally altered and mineralized to varying degrees 

(Bove and others, this volume). A regional propylitic alteration 

affects nearly all rocks in the study area (Burbank and Luedke, 

1969; Bove and others, this volume). The propylitic rocks 

are characterized by varying amounts of chlorite, epidote, 

calcite, and illite, in the presence of fresh to weakly altered 

primary feldspar crystals. Sparse, very fine grained pyrite also 

is common in the alteration assemblage. These weakly altered 

rocks were subjected to several generations and types of more 

intense alteration and mineralization. Studies by Bove and 

others (this volume) have delineated five different areas of the 

watershed that represent these different styles and ages of mineralization and alteration. Sampling locations and ranges of 

δ34S for the dissolved sulfate are shown relative to these areas 

in figure 1. Some of these locations fall outside the five area 

boundaries and are generally associated with characteristics 

of both the Eureka Graben and South Silverton mineralized 

areas. A few others have similarities to Red Mountain but also 
Figure 1. Animas River watershed study area, showing sampling site locations and distribution of δ34S values of dissolved sulfate in surface 

water. Extent of five outlined mineralized areas defined by Bove and others (this volume, Chapter E3). Note that “Mount Moly” was used by Bove and others to name a mineralized area but is not a formal geographic name.

EUREKA GRABEN

SOUTH

SILVERTON

OHIO PEAK-ANVIL 

MOUNTAIN

RED MOUNTAINS

MOUNT MOLY

34S of dissolved

 sulfate, per mil

-8 to -4 -4 to +0 

+0 to +4+4 to +11

Cement Creek

An mi as River

South 

Fork

Middle Fork

Base from U.S. Geological Survey1:100,000 Digital Line GraphsAlbers Equal-Area Conic projectionStandard parallels 37o30' and 40o30'

Central meridian 37o00

Watershed studyarea boundary

Mineralized area

107O37'30" 107O45'

37O

52'30"

37O

45'

0

0

1

1

2

2

3

3

4 MILES

4 KILOMETERS 

Mineral Creek

EXPLANATION

398 Environmental Effects of Historical Mining, Animas River Watershed, Colorado

'
Aqueous-Sulfate Stable Isotopes and Acidic Drainage 399

contain other overlapping mineralization. The five mineralized 

areas, which include Mount Moly1

, Red Mountains, Ohio 

Peak–Anvil Mountain, Eureka Graben, and South Silverton, 

are discussed briefly in the following section. For a more 

detailed discussion of these mineralized areas, the reader is 

referred to Bove and others (this volume). Aqueous sulfate 

in different types of altered areas shows different ranges and 

average isotopic compositions in figure 1. The sulfur-bearing 

mineralogy of these mineralized and altered areas varies, as 

does the local hydrology, and together these aspects undoubtedly affect the fine-scale isotopic composition of aqueous 

sulfate in a given drainage.

Hydrothermal alteration in the Mount Moly area was 

associated with quartz-monzonite intrusive activity that produced subeconomic copper-molybdenum porphyry mineralization (Bove and others, this volume). In the center of this 

area is a zone of quartz-sericite-pyrite (QSP) altered rock 

that grades outward into weak sericite-pyrite and propylitic 

zones. Rocks in the QSP zone are composed of coarse-grained 

sericite, secondary quartz, and fine-grained pyrite (2–5 volume 

percent). Drill core data indicate that gypsum is present in 

thin veins and fractures below the oxidized surface of the 

QSP zone. Mineralization in the Mount Moly area produced 

quartz-molybdenite stockwork veins, which are present in the 

QSP zone, and pyrite-quartz veins, which are present on the 

margins of the hydrothermal system. Barite is found as vein 

fillings, and coarse-grained gypsum also is found in the pyritequartz veins.

Intrusion of high-level dacite porphyry in the Red 

Mountains and Ohio Peak–Anvil Mountain areas resulted 

in extensive zones of acid-sulfate alteration (Burbank and 

Luedke, 1969; Bove and others, this volume). The acid-sulfate 

assemblage is composed of quartz, alunite, pyrophyllite, dickite, and as much as 30 volume percent pyrite. Broad expanses 

of QSP-altered rocks also are present between the acid-sulfate 

centers. Abundant veins and veinlets of anhydrite are present at depth within the QSP and potassic assemblages but are 

absent within the acid-sulfate zones. In the Red Mountains 

area, mineral deposits consist of silver-copper-lead-arsenic 

ores hosted within breccia pipes and brecciated fault zones 

(Burbank and Luedke, 1969; Bove and others, this volume). 

By contrast, the Ohio Peak–Anvil Mountain system is largely 

devoid of significant mineral deposits.

The Eureka Graben and South Silverton areas are characterized by 18–10 Ma vein mineralization that appears to be 

spatially and perhaps genetically related to intrusions of highsilica alkali rhyolite (Bove and others, this volume). Veins 

are mostly a polymetallic variety (silver, lead, zinc, copper, 

±gold) and formed as fracture or fissure-fillings in the propylitic country rock. Many of these veins in the Eureka graben 

area are especially rich in manganese silicate gangue minerals. 

Vein and fracture-filling anhydrite and gypsum are also associated with the manganese-rich ores. Selenite (large clear blades 

of gypsum) after anhydrite is common in fractures throughout 

the watershed. Zones of hydrothermally altered rock related to 

these veins consist of narrow envelopes that are superimposed 

over regionally propylitic-altered rock. In some areas with 

high vein concentrations, the associated alteration zones can 

be locally pervasive.

This brief geologic overview provides a context for two 

important aspects of interpreting the stable sulfate isotope 

data: the mineralogy or source material and the spatial distribution of that source material. Sulfides, especially pyrite, are 

prominent throughout the Animas River watershed study area, 

and gypsum and anhydrite are also prominent in many places. 

Both of these sources contribute to the dissolved sulfate concentration of the surface and ground water. The remainder of 

this report focuses on the collection, measurement, and interpretation of the sulfate stable isotope data in the study area.

Methods

Water-Quality Sampling Methods

Water-quality samples for dissolved constituent and 

stable isotope analyses were collected at more than 80 

stream, spring, and inactive mine sites in the study area during 

low-flow conditions from 1994 to 1999. Most samples were 

collected in the Middle Fork Mineral Creek subbasin of the 

Mineral Creek basin, and the Prospect Gulch, Ohio Gulch, 

South Fork Cement Creek, and Topeka Gulch subbasins of 

the Cement Creek basin (table 1; fig. 1). Field measurements, 

including stream discharge, specific conductance, pH, and 

water temperature, were made at sampling sites using methods described by Wilde and Radtke (1998). Field meters were 

calibrated in the morning on the day of the sample collection. 

The pH meter was calibrated with pH 2 and pH 4 buffers 

or pH 4 and pH 7 buffers to bracket the range of pH values 

measured in the field. The specific-conductance (SC) meter 

was calibrated with two standards that bracketed the range of 

values expected at the sampling sites. Instantaneous discharge 

was measured using a current meter according to methods 

described by Rantz and others (1982). Samples for dissolved 

constituents were field filtered into 250 mL polyethylene 

bottles using disposable 0.45 μm cartridge filters. Sample aliquots for calcium and total-iron analyses were preserved with 

1 mL of concentrated nitric acid, and aliquots for ferrous iron 

analyses were preserved with 1 mL of 6M HCl. Unpreserved 

filtered aliquots were collected for sulfate determination. 

Unfiltered samples for stable isotopes of water were collected 

in 60 mL glass bottles with polyseal caps. Samples for sulfur 

and oxygen isotopes of sulfate were filtered into 1 L polyethylene bottles and preserved with 1 mL of 6M HCl.

1

Mount Moly is the informal name for the peak referred to in this volume 

as peak 3,792 m. Its use in this report is confined to repetition of the name 

given by Bove and others (this volume, Chapter E3) to the altered and mineralized area around that peak and does not indicate a formal geographic name 

for peak 3,792 m.
Site 

Number Description Date Sample type Cat Q cfs Field Spec. Cond. μS/cm Field pH Water temp. °C

CC1 Upper spring, Cement Creek 4/25/97 Spring 1 0.22 1,266 3.06 21.4

CC126 Spring in Topeka Gulch 9/2/94 Spring 1 0.030 835 3.77 7.4

CC127 Tributary of Topeka Gulch 9/2/94 Stream 1 0.040 642 3.56 12.4

CC128 Spring in Topeka Gulch 9/3/94 Spring 1 0.010 975 2.90 6.7

CC129 Topeka Gulch upstream from unnamed mine 9/3/94 Stream 1 0.15 558 3.27 9.0

CC130 Ferricrete spring in Topeka Gulch 9/4/94 Spring 1 0.001 437 4.32 7.5

CC131 Tributary of Topeka Gulch 9/4/94 Stream 1 0.002 575 3.49 17.6

CC132 Topeka Gulch upstream from MS73 9/9/94 Stream 1 0.030 500 3.45 10.3

CC32 South Fork Cement Creek downstream from 

Velocity Lake

10/17/96 Stream 1 0.43 150 7.10 3.0

CC36 Spring downstream from Red Mountain # 3 

in Prospect Gulch

9/8/97 Spring 1 0.005 654 3.25 1.1

CC44 Spring in Prospect Gulch 9/9/97 Spring 1 0.020 685 6.40 1.6

CC45 Spring in Prospect Gulch 9/9/97 Spring 1 0.012 78 3.80 7.8

CC50 Spring in Ross Basin 9/5/97 Spring 1 0.001 87 6.92 2.2

CC51 Spring in Ross Basin 9/5/97 Spring 1 0.010 287 6.96 1.2

CC52 Spring in South Fork Cement Creek 10/18/96 Spring 1 0.031 520 3.95 1.8

CC53 Spring near Corkscrew Pass 9/7/97 Spring 1 0.010 76 3.92 11.2

CC70 Stream in Prospect Gulch 9/10/97 Stream 1 0.055 158 6.83 10.0

CC86 Upper Georgia Gulch 9/25/97 Stream 1 0.282 316 6.54 9.2

MC38 Spring southwest of Ohio Peak 8/9/99 Stream 1 -- 1,340 2.58 --

MC61 Moss Bog spring near peak 3,792 m 8/26/97 Spring 1 -- 304 4.20 12.3

MC62 Iron Bog spring near peak 3,792 m 8/26/97 Spring 1 -- 659 3.45 8.5

MC69 Upper Red tributary 9/19/95 Stream 1 1.1 298 4.92 6.7

MC71 Spring in Red tributary 9/19/95 Spring 1 0.001 1,480 3.06 6.3

MC72 Small stream in Red tributary 9/19/95 Stream 1 0.17 881 3.12 13.4

MC74 Red tributary at mouth 9/20/95 Stream 1 1.6 1,087 3.39 3.8

MC76 Crystal Lake outflow 9/18/95 Stream 1 0.50 60 4.47 9.9

MC77 Spring near Ophir Pass 9/18/95 Spring 1 0.020 156 5.32 1.8

MC81 Moraine spring in Paradise basin 10/11/95 Spring 1 0.34 451 6.84 1.8

MC82 Spring in upper Paradise basin 10/11/95 Spring 1 0.070 1,845 5.72 5.5

MC83 Sulfide spring in Paradise basin 10/11/95 Spring 1 0.050 1,630 5.43 6.6

UA40 Spring near Houghton Mountain 9/4/98 Spring 1 0.019 609 3.95 0.6

UA40 Spring near Houghton Mountain 9/4/98 Spring 1 0.019 609 3.95 0.6

CC29 Red spring, lower Prospect Gulch 2/12/97 Spring 2 0.28 707 3.01 3.8

CC34 Spring above Big Colorado mine 10/18/96 Spring 2 0.46 295 3.75 4.1

CC57 Stream in Prospect Gulch 9/9/97 Stream 2 0.008 608 2.98 8.6

CC71 Stream in Prospect Gulch 9/26/97 Stream 2 0.18 365 6.90 4.8

MC30 Battleship Slide spring near peak 3,792 m 8/28/97 Spring 2 0.001 576 3.02 15.8

MC85 Spring in Middle Fork Mineral Creek 9/19/95 Spring 2 0.020 182 6.83 8.7

MC86 Spring in Middle Fork Mineral Creek 9/14/95 Spring 2 0.020 106 6.56 4.3

CC40 Spring in Middle Fork Cement Creek 10/1/97 Spring 3 0.045 195 3.77 3.8

CC69 Stream in Prospect Gulch 9/10/97 Stream 3 0.020 428 3.50 8.0

A68 Animas River at Silverton 2/25/97 Stream 4 18 360 6.90 0.1

A68 Animas River at Silverton 2/25/97 Stream 4 18 360 6.90 0.1

A72 Animas River downstream from Silverton 2/26/97 Stream 4 65 642 4.90 0.7

A72 Animas River downstream from Silverton 2/26/97 Stream 4 65 642 4.90 0.7

C48 Cement Creek at Silverton 9/29/94 Stream 4 17 1,040 3.94 9.6

CC100 Seep below Yukon tunnel dump 5/15/97 Spring 4 -- 5,480 2.57 10.4

CC133 Topeka Gulch downstream from MS73 9/9/94 Stream 4 0.17 1,530 7.22 7.6

Table 1. Chemical and isotopic data for non-mining affected and mining-affected samples collected in the Animas River watershed 

study area.

[Cat, category 1, no evidence of mining activity at site; 2, a site that appears to be unaffected, but is not unequivocal; 3, a site that is not directly impacted, 

but where upgradient mining activity likely affected the water quality; 4, direct discharge from mining activity at site. Calcium, sulfate, ferrous and ferric iron 

concentrations given in milligrams per liter, mg/L; --, no data; Q, discharge in cubic ft per second, cfs; Field Spec. Cond., specific conductance measured in the 

field, expressed in microsiemens per centimeter, μS/cm]

400 Environmental Effects of Historical Mining, Animas River Watershed, Colorado
Aqueous-Sulfate Stable Isotopes and Acidic Drainage 401

d18OH2O dDH2O d18OSO4 d34SSO4

Calcium 

mg/L

Sulfate 

mg/L

Molar 

Ca/SO4

 ratio

Ferrous iron 

mg/L

Total iron 

mg/L

Percent charge 

balance

–16.1 –117 0.90 6.80 160 634 0.61 -- 38.5 –15.3

–16.3 –117 –4.40 –3.20 75 480 0.37 38.6 45.0 –10.0

–15.7 –114 –4.70 –1.50 56 370 0.36 25.5 31.0 –11.7

–16.2 –118 –6.20 –3.80 48 400 0.29 0.80 13.0 –9.3

–15.4 –111 –5.90 –5.10 31 190 0.39 2.9 13.0 6.5

–16.5 –121 –3.90 –6.60 69 290 0.57 20.0 20.0 0.4

–15.6 –116 –3.70 –6.70 55 230 0.57 0.10 0.67 –2.7

–15.0 –110 –6.80 –4.80 38 190 0.48 1.5 3.6 –0.8

-- -- –3.10 –1.50 22 33 1.64 0.006 <0.03 27.0

-- -- –5.60 –4.10 35 238 0.35 0.87 5.2 4.7

–15.8 –112 –7.80 –1.90 111 320 0.83 0.012 0.048 –24.6

–16.4 –121 –6.40 –4.90 6 32 0.47 0.49 0.69 9.6

-- -- -- 0.00 13 28 1.10 -- <0.030 –8.2

-- -- -- –0.40 42 113 0.89 -- <0.030 –3.9

-- -- –2.30 –0.20 88 175 1.20 0.043 0.073 46.0

-- -- -- –4.60 1 25 0.14 -- <0.030 –43.3

–14.9 –107 –5.40 –0.60 23 72 0.79 0.11 0.064 –10.5

-- -- -- –3.20 44 143 0.74 0.03 <0.030 –10.6

-- -- -- –2.40 71 568 0.30 -- 52.0 –22.5

–15.9 –115 1.00 8.20 42 134 0.75 0.37 0.42 –6.2

–16.3 –118 –1.80 2.30 30 323 0.23 64.0 63.9 15.8

–16.2 –115 0.00 4.80 43 130 0.79 0.82 0.87 –0.7

–15.1 –107 2.20 8.00 280 780 0.86 0.22 1.3 8.4

–15.8 –113 –2.50 1.70 20 460 0.10 11.6 22.4 6.5

–15.9 –115 –1.20 3.20 99 660 0.36 49.3 58.3 11.4

–16.1 –116 –5.00 –2.20 4 28 0.30 0.25 0.32 –25.5

–14.1 –98 –7.60 –7.60 20 61 0.79 <.001 <0.001 7.3

–15.8 –111 0.50 7.10 82 210 0.94 <0.001 <0.001 –2.6

–15.2 –107 1.00 7.00 420 1,200 0.84 8.8 9.0 –11.2

–13.8 –101 2.30 7.10 360 1,000 0.86 0.42 0.43 –6.4

-- -- -- –2.60 22 363 0.15 0.01 <0.030 –14.4

-- -- -- –2.60 22 363 0.15 0.01 <0.030 –14.4

–16.5 –120 –3.70 –0.50 31 308 0.24 -- 45.6 –13.9

–16.0 –115 –3.60 0.90 26 89 0.69 0.050 0.083 12.2

–10.60 –2.80 17 145 0.28 0.41 8.2 –5.3

–15.2 –109 –7.00 –1.80 60 153 0.94 0.12 0.068 0.9

–14.4 –104 –3.40 3.60 22 182 0.29 0.14 4.3 3.8

–16.6 –121 –6.80 –0.80 29 54 1.29 <0.001 <0.001 7.3

–17.0 –123 –8.50 –2.80 17 25 1.63 <0.001 <0.001 4.2

–13.7 –95 –9.60 –2.40 12 61 0.45 0.023 0.24 –1.1

–15.4 –111 –6.60 –3.70 48 195 0.58 0.033 0.44 –8.3

–15.7 –113 –4.20 4.30 60 141 1.02 -- 0.079 2.2

–15.7 –113 –4.20 4.30 60 141 1.02 -- 0.079 2.2

–15.7 –115 –2.10 6.50 113 285 0.95 -- 2.7 12.9

–15.7 –115 –2.10 6.50 113 285 0.95 -- 2.7 12.9

–16.0 –116 –3.40 7.30 190 560 0.81 4.2 3.9 –4.4

-- -- –7.7 0.8 454 2,720 0.40 -- 1,250 --

–16.1 –118 0.30 10.00 380 930 0.98 7.6 8.1 2.7

Table 1. Chemical and isotopic data for non-mining affected and mining-affected samples collected in the Animas River watershed 

study area.—Continued
Site 

Number Description Date Sample type Cat Q cfs Field Spec. Cond. μS/cm Field pH Water temp. °C

CC134 Topeka Gulch at mouth 9/9/94 Stream 4 0.19 1,480 6.72 10.1

CC135 Tributary of Ohio Gulch 9/29/94 Stream 4 0.001 980 3.34 9.4

CC22 Prospect Gulch upstream from mouth 11/5/97 Stream 4 0.440 415 3.23 0.1

CC26 Prospect Gulch near mouth 9/8/97 Stream 4 0.56 526 3.22 8.9

CC43 Stream in Prospect Gulch 9/9/97 Stream 4 0.010 370 2.99 9.9

M27 Mineral Creek upstream from South Fork 10/5/95 Stream 4 27 542 4.99 4.4

M34 Mineral Creek at Silverton 10/5/95 Stream 4 48 359 6.90 7.8

MC79 Middle Fork Cement Creek near mouth 9/27/95 Stream 4 7.9 796 4.75 4.7

MC80 Middle Fork Mineral Creek upstream 

from Bonner mine

9/30/95 Stream 4 10 727 4.89 5.7

MC88 Middle Fork Mineral Creek upstream 

from Red tributary

9/20/95 Stream 4 5.6 778 6.39 6.7

MC89 Middle Fork Mineral Creek downstream 

from Red tributary

9/20/95 Stream 4 7.2 800 4.58 8.0

MC91 Spring below Independence mine 9/26/95 Spring 4 0.020 594 3.45 2.9

MC92 Spring below Bonner mine 9/26/95 Spring 4 0.15 970 3.12 3.9

MS13 Eveline mine 2/13/97 Adit 4 0.014 493 2.81 3.1

MS15 Forest Queen mine 2/12/97 Adit 4 0.036 883 4.44 7.8

MS15 Forest Queen mine 2/12/97 Adit 4 0.036 883 4.44 7.8

MS19 Joe and Johns mine 9/8/97 Adit 4 0.006 1,087 2.79 9.9

MS35 Unnamed mine in Prospect Gulch 9/25/97 Adit 4 0.002 420 6.86 3.0

MS5 Big Colorado mine 10/18/96 Adit 4 0.016 660 3.17 6.6

MS52 Silver Ledge mine 10/18/96 Adit 4 0.75 1,090 5.70 5.6

MS56 Yukon tunnel 5/15/97 Adit 4 -- 956 6.98 14.5

MS59 Galena Queen shaft 9/9/97 Adit 4 -- 1,062 2.60 4.2

MS6 Black Hawk mine 10/1/97 Adit 4 0.60 1,020 6.95 7.8

MS60 Hercules shaft 9/9/97 Adit 4 -- 1,780 2.35 5.8

MS61 May Day mine 5/15/97 Adit 4 0.030 1,810 2.64 1.1

MS69 Leib’s hand auger hole #1 5/15/97 Spring 4 0.100 1,123 2.58 4.4

MS72 Unnamed mine in Topeka Gulch 9/8/94 Adit 4 0.009 1,000 6.95 5.5

MS73 Unnamed mine in Topeka Gulch 9/9/94 Adit 4 0.14 1,720 6.85 5.9

MS74 Unnamed mine in Ohio Gulch 9/28/94 Adit 4 0.001 1,110 2.84 8.2

MS75 Unnamed mine in Ohio Gulch 9/28/94 Adit 4 0.026 1,540 3.57 6.0

MS76 Ruby Trust mine 9/18/95 Adit 4 1.0 560 6.36 7.1

MS77 Paradise Portal 9/28/95 Adit 4 0.60 1,962 5.70 4.6

MS78 unnamed mine in Middle Fork 

Mineral Creek

9/19/95 Adit 4 0.030 215 5.54 5.3

MS80 Independence mine 9/26/95 Adit 4 0.040 667 3.19 4.0

MS81 Koehler tunnel 9/25/95 Adit 4 0.020 3,520 2.45 2.3

MS82 Junction mine 9/25/95 Adit 4 0.15 3,310 2.50 2.5

MS83 Old Hundred mine in tunnel 8/24/95 Adit 4 0.50 570 6.82 10.2

MS83 Old Hundred mine in tunnel 8/24/95 Adit 4 0.50 570 6.82 10.2

MS84 Old Hundred mine at portal 8/24/95 Adit 4 1.5 590 7.35 15.0

MS84 Old Hundred mine at portal 8/24/95 Adit 4 1.5 590 7.35 15.0

MS85 Klondike mine in tunnel 8/25/95 Adit 4 0.001 500 7.58 1.9

MS85 Klondike mine in tunnel 8/25/95 Adit 4 0.001 500 7.58 1.9

MS86 Klondike mine at portal 8/25/95 Adit 4 0.020 123 6.32 2.7

MS86 Klondike mine at portal 8/25/95 Adit 4 0.020 123 6.32 2.7

MS87 American tunnel inside tunnel 8/22/95 Adit 4 -- 3,050 6.16 10.6

MS88 American tunnel at portal 8/22/95 Adit 4 -- 2,010 5.98 11.7

Table 1. Chemical and isotopic data for non-mining affected and mining-affected samples collected in the Animas River watershed 

study area.—Continued

402 Environmental Effects of Historical Mining, Animas River Watershed, Colorado
Aqueous-Sulfate Stable Isotopes and Acidic Drainage 403

d18OH2O dDH2O d18OSO4 d34SSO4

Calcium 

mg/L

Sulfate 

mg/L

Molar 

Ca/SO4

 ratio

Ferrous iron 

mg/L

Total iron 

mg/L

Percent charge 

balance

–15.6 –116 –0.70 8.90 330 850 0.93 0.90 0.87 –1.0

–12.4 –97 –7.96 –2.60 110 490 0.54 0.001 2.3 –6.7

-- -- –9.4 –3.2 25.5 132 0.46 1.04 3.15 --

–15.6 –113 –5.70 –1.20 29 227 0.30 0.42 26.8 15.7

–15.1 –109 –9.20 –2.30 8 92 0.21 0.043 1.9 18.5

-- -- –0.60 4.90 180 270 1.60 -- 5.6 67.2

-- -- –1.50 5.20 130 160 1.95 -- 1.5 81.0

–16.0 –116 0.00 4.80 130 430 0.73 15.5 15.6 1.4

-- -- –0.40 4.90 -- 410 -- -- 12.0 --

–16.2 –116 0.20 5.40 140 430 0.78 9.9 10.4 –13.5

–16.1 –115 –0.70 4.80 130 440 0.71 16.5 18.1 –0.6

–15.8 –116 –0.60 5.60 63 260 0.58 0.58 2.4 –2.0

–16.1 –117 0.20 6.40 110 440 0.60 0.50 4.7 1.2

–16.8 –122 –6.60 –2.60 4 153 0.07 -- 14.7 30.2

-- -- –4.20 5.90 133 450 0.71 23.3 24 --

-- -- –4.20 5.90 133 450 0.71 23.3 24 --

–17.0 –123 –7.60 –4.40 2 367 0.01 0.91 60.3 –10.8

–16.2 –117 –6.90 –0.20 79 143 1.32 0.012 <0.03 1.7

–16.1 –116 –6.60 –2.60 64 277 0.55 0.40 1.4 5.0

–15.8 –114 –1.60 7.30 211 518 0.98 11.8 12.3 8.0

–15.5 –111 –0.70 9.70 178 538 0.79 -- 0.22 –30.9

–16.7 –121 –7.60 –1.70 7 368 0.05 0.83 60.1 23.5

–14.9 –107 –4.30 6.60 259 600 1.04 0.80 1.2 5.4

–15.9 –115 –8.40 –3.10 3 691 0.01 1.8 118 3.2

–14.5 –113 –9.40 –4.90 27 1,060 0.06 -- 189 –61.0

-- -- –10.8 –3.4 9.32 300 0.07 -- 37.8 --

–15.8 –116 –1.20 8.00 340 770 1.06 4.0 3.8 0.3

–16.4 –120 0.83 10.60 460 1,100 1.00 11.7 12.0 2.7

–14.9 –109 –11.10 –3.40 26 520 0.12 0.004 78.0 –5.9

–15.9 –117 –4.80 –0.50 180 950 0.45 0.094 13.0 –35.1

–16.6 –121 –3.40 5.20 110 200 1.32 0.46 0.47 9.3

–16.0 –115 –0.50 4.70 400 1,200 0.80 66.5 67.0 8.6

–16.7 –122 –5.80 –1.40 20 83 0.58 9.1 9.4 –0.8

-- -- –0.10 5.80 65 -- -- 1.4 12.6 --

-- -- –10.70 –0.40 111 2,720 0.10 27.8 686 41.0

-- -- –9.50 –0.60 144 2,700 0.13 31.7 494 4.4

–15.4 –110 –6.60 –1.20 110 200 1.32 <0.001 0.005 6.4

–15.4 –110 –6.60 –1.20 110 200 1.32 <0.001 0.005 6.4

–15.4 –112 –3.20 4.00 110 210 1.26 <0.001 0.005 7.3

–15.4 –112 –3.20 4.00 110 210 1.26 <0.001 0.005 7.3

–14.6 –103 –5.40 –0.70 98 150 1.57 <0.001 0.006 8.3

–14.6 –103 –5.40 –0.70 98 150 1.57 <0.001 0.006 8.3

–15.4 –110 –5.40 –1.50 23 40 1.38 <0.001 0.005 10.4

–15.4 –110 –5.40 –1.50 23 40 1.38 <0.001 0.005 10.4

–16.3 –117 –4.90 1.90 510 2,100 0.58 114 310 –6.6

–16.1 –117 –2.60 7.80 470 1,300 0.87 27.3 27.0 1.9

Table 1. Chemical and isotopic data for non-mining affected and mining-affected samples collected in the Animas River watershed 

study area.—Continued
Analytical Methods for Dissolved Constituents

Samples collected in 1994–95 were analyzed at the 

USGS National Water Quality Laboratory in Colorado (URL 

http://wwwnwql.cr.usgs.gov/, accessed November 2002). Calcium and total iron were determined by inductively coupled 

plasma–atomic emission spectroscopy, and sulfate was determined by ion chromatography according to methods published 

by Fishman and others (1994). Samples collected in 1996–99 

were analyzed at a USGS research laboratory in Boulder, 

Colo. Calcium and total iron were determined by inductively 

coupled plasma–atomic emission spectroscopy, and sulfate 

was determined by a turbidimetric method using a Hach 

SulfoVer reagent and Hach DR2000 spectrophotometer. For all 

years, ferrous iron was determined by a colorimetric method 

using a Hach 1, 10-phenanthroline reagent and Hach DR2000 

spectrophotometer. The quality of laboratory analyses was 

assessed through the analysis of laboratory blanks, sample replicates, USGS standard reference water samples (URL http://

btdqs.usgs.gov/srs/, accessed November 2002), and calculated 

ion balances. Stable isotopes of water (δ18O and δD) were 

analyzed at the USGS stable isotope laboratory in Reston, Va. 

(URL http://isotopes.usgs.gov/index.htm/, accessed November 

2002).

Analytical Methods for Oxygen and Sulfur 

Isotopes of Dissolved Sulfate

One-liter samples collected for sulfate isotopes were 

pumped through two sequential columns using a peristaltic 

pump at a flow rate of 4 mL per minute. The first column 

was filled with a chelating resin (BioRad Chelex 100) and 

was used to remove iron from solution, because iron can 

act as an interferant during the isotope analyses (C. Janik, 

oral commun., 1994). The second column was filled with a 

NaCl-saturated anion exchange resin (BioRad AG 1-X8 anionexchange resin, 200–400 mesh), which was used to collect 

40 mg of sulfate from each sample. After the sample was 

pumped through both columns, sulfate was eluted from the 

anion exchange column using 0.4 M KCl, then precipitated as 

BaSO4

 by adding saturated BaCl2 to the eluent, and heated to 

enhance crystallization of the precipitate. The BaSO4

 precipitate was collected on a membrane filter and air dried prior to 

analysis. Oxygen and sulfur isotopic analyses were performed 

on the BaSO4

 precipitate by isotope-ratio mass spectrometry 

at the USGS stable isotope laboratory in Menlo Park, Calif., 

in 1994 and in Denver, Colo., from 1995 to 1999. Oxygen 

isotopes of dissolved sulfate are expressed in per mil (parts 

per thousand) relative to the Vienna Standard Mean Ocean 

Water (VSMOW) on a scale that is normalized such that 

the δ18O of the Standard Light Arctic Precipitation (SLAP) 

water is –55.5 per mil exactly (Coplen, 1996). The δ34S of 

dissolved sulfate are expressed in per mil relative to Cañon 

Diablo troilite (CDT) standard. Precision of the analyses is 

±0.2 per mil.

Analytical Methods for Oxygen and Sulfur 

Isotopes of Minerals

Pyrite, gypsum, and anhydrite were purified by hand 

picking. Gypsum and anhydrite were dissolved in distilled 

water, filtered, and the dissolved sulfate precipitated as 

BaSO4

 with addition of BaCl2 with heating after acidification 

with HCl to pH 3. Isotopic compositions of sulfur-bearing 

minerals were made by several methods over the course of 

the study. Initial sulfur isotope analyses of BaSO4

 precipitates and pyrite were made using conventional methods for 

the preparation of SO2

 gas (Yanagisawa and Sakai, 1983) 

with analyses on a 6-60 RMS Nuclide Corporation mass 

spectrometer with a modified source, collector, and pumping 

system. Later sulfur-isotope measurements were made on a 

Micromass Optima mass spectrometer using continuous flow 

methods similar to those described by Giesemann and others (1994). Initial oxygen-isotope analyses of sulfates were 

made on CO2

 prepared using a BrF5 technique described by 

Wasserman and others (1992). Later BaSO4

 samples were 

analyzed as CO by continuous flow methods using a hightemperature pyrolysis device on a Thermo Finnigan Delta 

Plus mass spectrometer.

Results and Discussion

Seventeen pyrite samples were analyzed for δ34S and 

4 gypsum and 2 anhydrite samples were analyzed for both 

δ34S and δ18OSO4

. Results of these analyses are shown in 

table 2. The distinct isotopic signatures between the pyrite 

and the gypsum/anhydrite samples can be clearly seen in 

figure 2. The δ34S of the pyrite ranges from +2.5 to –6.9 per 

mil, whereas those for gypsum and anhydrite samples vary 

from +15 to +18 per mil. This large separation in sulfur 

isotopic composition provides a useful tracer for estimating 

the amounts of gypsum/anhydrite dissolution and pyrite oxidation contributing to sulfate in water from the study area, 

but these results also confirm that the gypsum/anhydrite is 

of hypogene origin, and not from reaction of acid sulfate 

waters with calcite or plagioclase during weathering.

Ninety water samples were collected for δ34SSO4 determinations, and most of these were also analyzed 

for δ18OSO4

, δD, δ18OH2

O, discharge, specific conductance, 

pH, temperature, and inorganic constituents. The results used 

for interpretations in this report are compiled in table 1. For 

the purposes of this discussion, only the calcium, sulfate, and 

iron concentrations are reported here. The complete analyses 

of dissolved inorganic constituents for these samples are 

given in Sole and others (this volume, Chapter G).

404 Environmental Effects of Historical Mining, Animas River Watershed, Colorado
Aqueous-Sulfate Stable Isotopes and Acidic Drainage 405

Distribution of Stable Sulfate Isotopes for Mined 

and Unmined Areas

Figure 3 shows the distribution of δ34S and δ18O for dissolved sulfate along with some limited data on isotope compositions for gypsum/anhydrite and pyrite. The box suggested for 

the range of stable isotope composition for gypsum/anhydrite 

likely extends to the right (towards higher δ18O SO4), shown by 

the area outlined by the dashed line, because extrapolation of 

the mixing trends in the dissolved stable sulfate isotope data 

indicates it and because two gypsum samples of unknown 

origin plot in this area of the box. The range of δ18O for 

water in the upper Animas River basin is shown as a vertical band on the left side of the diagram and the δ18O for air 

(+23.5 per mil) is shown as the right-hand border of the figure. 

The trend in these data indicates that dissolved sulfate results 

from a mixture of sulfate from gypsum/anhydrite dissolution 

and from pyrite oxidation. Note too that as δ18OSO4

 values 

become lower (more dominated by pyrite oxidation), they also 

become closer to δ18OH2

O. This trend is consistent with the 

Table 2. Sulfur and oxygen isotopic data for sulfide and sulfate minerals in the study area.

Sample No. Mineralization area1 Mineral Mineralogical notes2 δ34S δ18OSO4

B1-1918AN American tunnel anhydrite Vein anhydrite, 1,918' below American tunnel 17.3 –3.3

B1-1900 American tunnel anhydrite Vein anhydrite, 1,900' below American tunnel 18.0 1.1 

B1-1685 American tunnel gypsum Coarse, clear vein gypsum, 1,685' below 

American tunnel.

16.7 –1.9

PPSO4-1 Mount Moly gypsum White, fine-granular, vein margin at Paradise portal 14.6 4.7

NSM-1 Near South Silverton gypsum Coarse, clear vein gypsum; adjacent to QSP margin 15.2 1.1

TG-1 Ohio Peak–Anvil Mountain gypsum Coarse, clear vein? gypsum; mine dump lower 

Topeka Gulch.

17 3.6

B1-1918PY American tunnel pyrite Fine-grained pyrite in vein with anhydrite, 1,918' 

below American tunnel.

–1.6 --

SDY3997 Mount Moly pyrite Fine blebs of 1–2 mm disseminated pyrite in 

QSP zone.

2.5 --

SDY05A97 Mount Moly pyrite Fine blebs of 1–2 mm disseminated pyrite in 

QSP zone.

0.7 --

SDPH01PY Mount Moly pyrite Medium-grained pyrite, fracture in hyd prop zone 0.9 --

SDY10B98 Mount Moly pyrite Medium-grained pyrite in qtz vein central QSP zone –3.1 --

BONNERPY Mount Moly pyrite Late-stage pyrite in 3 mm vein 2.0 --

SDB39 Ohio Peak–Anvil Mountain pyrite Finely disseminated pyrite, QSP zone, 

mid Topeka Gulch.

–2.3 --

SDB38 Ohio Peak–Anvil Mountain pyrite Finely disseminated pyrite, QSP zone, 

mid Topeka Gulch.

–5.4 --

IDB61B Red Mountains pyrite Finely disseminated pyrite, QSP zone 

Prospect Gulch.

–1.5 --

IDB40 Red Mountains pyrite Finely disseminated pyrite, QSP zone 

Prospect Gulch.

–1.7 --

IDB98 Red Mountains pyrite Finely disseminated pyrite, QSP zone 

Prospect Gulch.

–4.7 --

IDB04PY Red Mountains pyrite Finely disseminated pyrite, quartz-pyrophyllitealunite zone Red Mtn #3.–5.0 --

HMT-804-602P Red Mountains pyrite Finely disseminated pyrite, quartz-alunite zone 

602' beneath upper Lark mine.

–6.9 --

LS998PY Red Mountains pyrite Fine-grained vein pyrite, Lark mine stage 3b 

(see text).

–3.6 --

LS2BPY Red Mountains pyrite Fine-grained vein pyrite, Lark mine stage 3b 

(see text).

–3.9 --

GQWSTPY Red Mountains pyrite Fine-grained vein pyrite, Prospect Gulch, above 

Galena Queen mine, stage 3b (see text).

–2.8 --

GQ498PY Red Mountains pyrite Fine-grained vein pyrite, Galena Queen mine, 

stage 3b (see text).

–3.4 --

1

Mount Moly, Ohio Peak-Anvil Mountain, and Red Mountains are names applied to mineralized and altered areas in Bove and others, this volume, 

Chapter E3. Mount Moly is an informal name applied to the area of peak 3,792 m between Middle and South Forks Mineral Creek.

2

QSP, quartz-sericite-pyrite; hyd prop, hydrothermal propylitic alteration; qtz, quartz. 
Figure 2. δ34S values for pyrite and gypsum/anhydrite from Animas River watershed 

study area.

0

1

2

3

4

5

6

7

-10 -5 0 5 10 15 20 25

NUMBER OF SAMPLES

Gypsum/anhydrite

Pyrite

34S, IN PER MIL

hypothesis that during pyrite oxidation most of the oxygen in 

the sulfate comes from water. Indeed, the plot deviates significantly from the diagonal mixing trend for the lowest isotope 

values where the values trend almost horizontally towards 

the δ18OH2

O values. There is also a much greater dispersion in 

the δ18OSO4 among the pyrite oxidation-dominated samples, 

and these isotopically light samples show noticeable overlap 

between unmined and mining-influenced sites.

Mining-influenced samples are differentiated from those 

that are not (or are minimally) influenced by mining activities in figure 3, according to a set of established criteria (Mast 

and others, this volume). A tendency for separation of δ18OSO4 values from mined compared to unmined environments can 

be seen, with the samples from unmined areas having higher 

values. (Although one could interpret the trend alternately 

in terms of the δ34SSO4

 values being lower, there is no known 

process by which this pattern could be explained.) The 

separation at heavier δ34SSO4 values is likely caused by differences in the isotopic composition of the predominant gypsum/

anhydrite dissolving in the basin rather than by different 

processes of weathering in mined and unmined environments. 

The tendency for separation among isotopically light samples 

is not as clear because of notable overlap but is, nevertheless, 

apparent. A possible reason for this tendency of the isotopic 

values to separate in the pyrite oxidation-dominated part of 

the diagram might be the inclusion of more molecular oxygen 

reacting with the pyrite in the unmined settings compared 

to the mined settings. However, air has more accessibility 

in the mined subsurface than in unmined aquifers—and the 

trend should be in the opposite direction. Qureshi (1986, Ph. 

D. thesis, University of Waterloo, Waterloo, Ontario, as cited 

in van Stempvoort and Krouse, 1994) found that as the PO2 increased during laboratory oxidation of pyrite, the resulting 

sulfate decreased in δ18O. He ascribed this trend to changes 

in the oxidative pathway and changing importance of intermediate sulfoxyanions during oxidation. Van Stempvoort and 

Krouse’s (1994) comment on Qureshi’s results considered 

the possibility that increasing PO2 may decrease sulfite-water 

oxygen isotope exchange because the half-life of sulfite would 

decrease. Decreasing exchange, however, would tend to drive 

the isotopic composition to heavier values, and consequently, 

the samples from mined environments should be heavier in 

isotopic composition than those from the unmined environments. The actual trend (fig. 3) is the opposite from this 

concept. Alternatively, if the assumption is made that microbially catalyzed oxidation is enhanced in mined environments 

and, therefore, a much greater concentration of ferric iron is 

produced in mines, we could conclude that amounts of water 

oxygen in the sulfate from mined environments should be 

greater because high ferric iron concentrations use water oxygen in the oxidative pathways.

Although δ18O values of aqueous sulfate in acidic drainage 

associated with unmined areas appear to be offset to heavier 

values relative to mined areas, spatial patterns (combined 

effects of variations in local geology, mineralogy, elevation, 

and local hydrology) in isotopic values must also be considered. 

For example, the Mount Moly alteration area (Bove and others, 

this volume) has acidic drainage largely unaffected by mining, 

but it does include a few mines whose water contains δ18OSO4 values nearly the same as those from the unaffected drainage (fig. 4). If these data were examined alone, the conclusion 

would be that there is no noticeable isotopic discrimination 

between drainage from unmined and from mined areas. The 

effect of spatial differences can be seen better in figure 4 where 

the data points from figure 3 are shown on an enlarged scale 

and with different symbols designating the different alteration areas (fig. 1) that have been sampled. Most of the data for 

both environments in the Mount Moly alteration area group 

closely together but separately from the Eureka Graben and 

South Silverton areas. These two distinct but parallel trends 

suggest that the δ18OSO4

 from dissolving gypsum and anhydrite 

from the Mount Moly area is just a few per mil heavierthan 

406 Environmental Effects of Historical Mining, Animas River Watershed, Colorado
Aqueous-Sulfate Stable Isotopes and Acidic Drainage 407

that typical of the other areas. One sample of selenite has been 

found in this area that has higher δ34SSO4 and δ18OSO4 values than 

do other gypsum samples.

Data for samples from the unmined and mined areas 

in Topeka and Ohio Gulches (Ohio Peak–Anvil Mountain 

area) and in the Eureka graben show the best tendency for 

this separation, and the data for the Prospect Gulch area 

(Red Mountains area) and the Mount Moly area show the least 

tendency (or the most overlap). There may be differences in 

the stable sulfate isotope signatures between undisturbed and 

mining-influenced drainage, but careful discrimination of the 

data must be made to properly address this trend based on 

several factors, including differences of the source gypsum/

anhydrite δ18OSO4

 values, pathways of water flow during 

weathering, and possible effects of pH and aqueous iron redox 

concentrations. Trends in sulfur and oxygen stable isotopes 

may follow the same pathways from an intrusive center as 

changes in mineral zonation and alteration. For example, gypsum and anhydrite are found throughout the watershed except 

in acid-sulfate alteration zones. Our examination of the data,

with regards to zoning, does not warrant further interpretation 

at this time.

Influence of Gypsum and Anhydrite on 

Water Chemistry

The influence of gypsum and anhydrite dissolution on 

water chemistry is shown by figure 5, a plot of calcium concentrations against sulfate concentrations. All samples with 

δ34SSO4

 greater than +4 per mil are shown by solid circles and 

those with values less than +4 per mil by open circles. The 

value of +4 per mil is used as a rough estimate for dividing 

samples dominated by sulfate from gypsum/anhydrite dissolution versus those dominated by sulfate from pyrite oxidation. 

Samples with higher sulfur isotope values plot near or on the 

dissolution line for congruent dissolution of pure gypsum or 

anhydrite. Indeed, some of the samples closely approach the 

stoichiometric solubility equilibrium value for gypsum saturation (the stable phase of solid calcium sulfate at temperatures 

below 56°C). Most of the samples with low calcium:sulfate 

ratios also have δ34SSO4

 less than +4 per mil, indicating that 

most of the aqueous sulfate is derived from the only other viable source, pyrite oxidation. If low calcium:sulfate ratios are 

indicative of pyrite oxidation, then the waters with the lowest 

Figure 3. δ34S of sulfate plotted against δ18O of sulfate for surface waters in Animas River watershed study 

area. Open symbols, background water samples; solid symbols, mining-affected water samples; dotted 

open symbols, gypsum/anhydrite samples. Dashed line, likely extension of range (text, p. 405).

-10

-5

0

5

10

15

-20 -10 0 10 20

 

18O OF SULFATE, IN PER MIL

34S OF SULFATE, IN PER MIL

Range of 34S 

for pyrite 

18O of air (+ 23.5 per mil) 

18O 

of H2O Range of 

34S for 

gypsum & anhydrite
ratios should have the lowest pH values. Figure 6 is a plot of 

the Ca/SO4

 molar ratio against pH with a solid horizontal line 

at 1.0 showing ideal gypsum/anhydrite congruent dissolution 

and an envelope of two dashed lines showing the range of 

gypsum/anhydrite-dominated dissolution from 0.6 to 1.1. The 

striking aspect of this plot is that those samples with low Ca/

SO4

 ratios (and low δ34SSO4) all occur at low pH values, nearly 

all below pH values of 4. The inferred gypsum-dominated 

samples range in pH from 3 to 7, and those with Ca/SO4 ratios 

higher than 1.1 are predominantly at pH values higher than 6, 

suggesting that calcite dissolution is as important as gypsum 

dissolution.

Other sulfate minerals known to occur in the watershed 

include barite, alunite, and jarosite, and these have some 

potential for contributing to the sulfate concentrations in the 

water during weathering. These sulfate minerals, however, and 

most others that may occur are much less soluble than gypsum 

and anhydrite. Such a large contrast in solubility means that 

the solubility behavior of barite, alunite, and jarosite will be 

determined by the concentrations of sulfate produced by the 

weathering rates of gypsum, anhydrite, and pyrite.

Three Dominant Processes Controlling 

Water Quality

The preceding trends indicate that three main water types 

can be classified in the Animas River watershed study area 

using Ca/SO4

 ratios and δ34SSO4 values, as shown in figure 7. 

Each end-member water type represents a major process 

dominating the water chemistry, and the drawn circles in figure 7 approximately outline these three end-member compositions. The remaining data points between the circles are mixtures between end members. Some of the data points within 

Figure 4. δ34S of sulfate plotted against δ18O of sulfate for surface waters in Animas River watershed 

study area. Symbols designate different alteration areas shown in figure 1. Open symbols represent 

unmined environments; solid symbols represent mined environments. OPAM, Ohio Peak–Anvil Mountain.

-10

-5

0

5

10

-12 -7 -2 3

Eureka Graben (unmined)

Alteration areas (fig. 1)

Eureka Graben (mined)

Mount Moly (unmined)

Mount Moly (mined)

OPAM (unmined)

OPAM (mined)

Red Mountains (unmined)

Red Mountains (mined)

34S OF SULFATE, IN PER MIL

 

18O OF SULFATE, IN PER MIL

EXPLANATION

408 Environmental Effects of Historical Mining, Animas River Watershed, Colorado
Aqueous-Sulfate Stable Isotopes and Acidic Drainage 409

0 10 20 30

SULFATE, IN MILLIMOLES PER LITER

0

5

10

15

20

CALCIUM, IN MILLIMOLES PER LITER

GYPSUM DISSOLUTION LINE

GYPSUM STOICHIOMETRIC SATURATION

0 2 4 6 8 10

pH, IN STANDARD UNITS

0

0.5

1.0

1.5

2.0

2.5

CALCIUM:SULFATE, MOLAR RATIO

Figure 5. Dissolved calcium versus dissolved sulfate, showing their distribution relative to 

the congruent dissolution line for pure gypsum (or anhydrite) and the equilibrium solubility 

value for gypsum in pure water. Solid symbols, water samples with δ34S greater than +4.0 per 

mil, indicating a predominance of dissolved gypsum/anhydrite; open symbols, δ34S <+4.0 per mil.

Figure 6. Molar ratio of calcium to sulfate plotted as a function of pH. Solid symbols, samples 

with δ34S values greater than +4.0 per mil, as in figure 5; open symbols, δ34S <+4.0 per mil. Solid 

horizontal line represents exact gypsum dissolution stoichiometry; two horizontal dashed lines 

represent approximate range of gypsum-dominated water compositions.
the drawn circles are also mixtures, and the circles are only 

shown to help delineate the compositions that are approximately dominated by one of the three processes. One cluster of 

values appears at low δ34SSO4

 and low Ca/SO4 ratios, indicative 

of acid drainage dominated by pyrite oxidation. Another cluster appears at Ca/SO4

 ratios close to 1.0 but with the heaviest 

δ34SSO4

 values, indicative of gypsum-dominated dissolution. A 

third, smaller cluster can be seen at intermediate δ34SSO4

 values 

between that of gypsum dissolution and pyrite oxidation but 

having the highest Ca/SO4 ratios, indicating the influence 

of calcite dissolution in addition to gypsum dissolution. For 

example, if calcium is solely from calcite and gypsum dissolution in a 1:1 mole ratio, then the Ca/SO4

 ratio would be 2. If 

the proportion of calcite dissolved was 50 percent of the molar 

quantity of gypsum dissolved, then the Ca/SO4

 ratio would 

be 1.5. Figure 7 provides the simplest overall picture of water 

quality demonstrating the three main processes dominating the 

water chemistry: (1) pyrite oxidation, (2) gypsum dissolution, 

and (3) calcite dissolution using only two variables (δ34SSO4 and Ca/SO4

 ratio).

The plots shown in figures 5, 6, and 7 imply or assume 

that dissolved aqueous calcium and sulfate behave conservatively; that is, the only reactions that can affect their concentrations are the dissolution of gypsum, anhydrite, and pyrite. 

Figure 5 shows that the conservative behavior assumption 

holds because of the consistent linear relationship between 

calcium and sulfate along the stoichiometric dissolution line 

for isotopically heavy samples. Although conservative behavior (in the strictest sense for all possible reactions) probably 

does not hold, the question is whether other side reactions 

such as the dissolution of plagioclase feldspar, the dissolution 

of alunite and jarosite, the precipitation of gypsum and calcite, 

the dissolution of other sulfide minerals (such as chalcopyrite, 

sphalerite, arsenopyrite), and the dissolution of pyroxenes and 

amphiboles have any significant effect on the overall mass 

balance derived from dissolution of gypsum, anhydrite, and 

pyrite. Based on relative solubilities, relative dissolution rates, 

relative abundances of minerals in the watershed, and the 

consistent patterns shown in figures 5–7, the assumption of 

conservative behavior for calcium and sulfate certainly seems 

justified in the overall mass balance.

Relationship Between δ18OSO4 from Pyrite 

Oxidation and δ18OH2O

From the previous discussion, the aqueous sulfate samples 

that are most dominated by pyrite oxidation are clearly those 

with the lowest pH, the lowest δ34SSO4 values, and the lowest Ca/SO4

 ratios. This characterization of the data allows the 

examination of pyrite-dominated samples for the relationship 

between the water oxygen and the sulfate oxygen. Figure 8 

plots all the data for δ18OSO4

as a function of δ18OH2

O, similar 

to the type of plot used by Taylor and Wheeler (1994) and 

van Stempvoort and Krouse (1994). The two diagonal lines are 

from the study by Gould and others (1989), in which pyrite 

and other sulfide minerals were allowed to oxidize in the 

presence of Acidithiobacillus ferrooxidans with increasingly 

Figure 7. δ34S of dissolved sulfate in streams, springs, and adits plotted against 

calcium:sulfate molar ratio.

-15

-10

-5

0

5

10

15

0 0.5 1.0 1.5 2.0

Streams

Springs

Adits

CALCIUM:SULFATE, MOLAR RATIO

34S OF SULFATE, IN PER MIL

Gypsum-dominated

Calcite-dominated

Pyrite-dominated

EXPLANATION

410 Environmental Effects of Historical Mining, Animas River Watershed, Colorado
Aqueous-Sulfate Stable Isotopes and Acidic Drainage 411

18O-enriched H2

O. Linear equations were derived from the 

data, and these lines represent the maximum and minimum 

values from that study. Our data show a range of values somewhat similar to those of other studies, but we have designated 

those samples with low values of δ34SSO4

 (same as those with 

low Ca/SO4 ratios less than 0.6) with open circles to differentiate against the gypsum-dominated samples (solid circles). 

These samples (open circles) should show the best correlation 

of δ18OSO4

with δ18OH2

O; however, they do not. Many of the 

samples do fall in the range found by Gould and others (1989) 

and Seal (2003), but a clear positive correlation is not evident 

in the data. Differentiating the more acidic samples (those 

with pH <3) does not help, although these do tend to have the 

lowest δ18OSO4

 values and they are closest to the δ18OH2

O values. 

The samples with the highest δ18OH2

O values are also not the 

most acidic samples. The tendency for some samples to have 

higher δ18OH2

O values suggests that evaporation (or mixing with 

evaporated waters) may have affected them. Evaporation occurring after sulfate formation would increase the δ18OH2

O without 

affecting the δ18OSO4

.

The water isotopes for these samples fall close to the 

Rocky Mountain meteoric water line (M.A. Mast, unpub. 

data, 2002) shown in figure 9. Figure 9B is an enlargement 

of the Animas River watershed study area samples and 

indicates that at least two samples were affected by evaporation. One of these two samples corresponds to the highest 

δ18OH2

O in figure 8 (tributary of Ohio Gulch having a low 

δ18OSO4

). The other sample is the May Day adit sample having 

a δ18OH2

O = –14.5 o

/

oo and a δ18OSO4

= –9.4 o/

oo, the third lowest 

in δ18OSO4shown in figure 8. The next most likely sample 

to have undergone evaporation is from a sulfide spring in 

Paradise Basin with δ18OH2

O = –13.8 o

/

oo and a δ18OSO4

= 2.3 o/

oo

shown as the highest δ18OH2

O for a solid circle in figure 8. 

Some of the other samples may be affected by evaporation 

to a lesser degree. Overall, no clear trend of δ18OSO4

 can be 

observed with δ18OH2

O for aqueous sulfate in samples dominated by pyrite oxidation unless the amount of evaporation 

and the amount of gypsum or anhydrite dissolution could 

be determined. Several other processes affect the isotopic 

composition of the aqueous sulfate from pyrite oxidation 

such that the δ18OSO4

 values are broadly dependent on δ18OH2

O

but not with respect to specific locations. Mixing of different 

waters, between the point of pyrite oxidation and the point of 

sampling, seems likely to change the δ18OH2

O without changing the δ18OSO4

. Similarly, dilution during snowmelt runoff and 

major rainstorms and changes in precipitation from one storm 

to another can change δ18OH2

O and δ2

H independently of other 

isotopes.

-18 -17 -16 -15 -14 -13 -12

-10

-5

0

EVAPORATION

 

18O OF H2

O, IN PER MIL

18O OF SULFATE, IN PER MIL

 

34S > +4 per mil

 

34S < +4 per mil, pH >3

 

34S < +4 per mil, pH <3

EXPLANATION

MIXING

MIXING

Figure 8. δ18O of sulfate plotted against δ18O of H2O, showing no dependence on water 

oxygen but possible dispersion of values at low pH due to evaporation. High values of 

δ18O of sulfate indicate influence from gypsum/anhydrite dissolution or mixing with such 

waters. Diagonal lines are based on experimental data from Gould and others (1989).
-125

-120

-115

-110

-105

-100

-95

-90

-18 -17 -16 -15 -14 -13 -12

D, IN PER MIL

 

18O, IN PER MIL

RMMWL 

 D = 8.1* 18O + 14.8

-200

-150

-100

-50

0

-25 -20 -15 -10 -5 0

Precipitation

Surface water

D, IN PER MIL

 

18O, IN PER MIL

 RMMWL

 D = 8.1* 18O + 14.8

GMWL

 D = 8.1* 18O + 10.8

EXPLANATION

A

B

Figure 9. δ18O versus δD of A, surface waters collected in Animas River watershed study area and precipitation 

collected in the Rocky Mountain region, shown with global meteoric water line (GMWL) and Rocky Mountain 

meteoric water line (RMMWL). B, Inset of upper graph, showing only surface water samples collected in study 

area, with RMMWL. GMWL is taken from Clark and Fritz (1997).

412 Environmental Effects of Historical Mining, Animas River Watershed, Colorado
Aqueous-Sulfate Stable Isotopes and Acidic Drainage 413

Conclusions

Analysis of more than 100 samples of waters and 

minerals for sulfur and oxygen isotopes in the Animas River 

watershed study area, coupled with interpretation, leads to 

several conclusions. Sulfur and oxygen isotope compositions 

of aqueous sulfate show a distinct mixing trend from those 

waters dominated by sulfate production from pyrite oxidation to those dominated by gypsum and (or) anhydrite dissolution. Some discrimination between non-mining affected 

acid-rock drainage and mining-influenced drainage is apparent in the data for two of the four areas studied in detail, but 

these differences may be more related to local compositions 

of primary hydrothermal sulfates than to geochemical processes related to weathering. In the study area, the isotopes 

show that the gypsum/anhydrite occurrences must be largely 

of hypogene (hydrothermal) origin, not of supergene (weathering) origin. The gypsum/anhydrite isotopic signature is most 

evident in samples with higher calcium and sulfate concentrations. Because of the nearly stoichiometric concentrations 

of calcium and sulfate in many samples on a molar basis, 

gypsum/anhydrite is apparently the primary source of calcium 

in these waters. Samples with low calcium:sulfate ratios have 

consistently low δ34SSO4

 values, indicating pyrite oxidation as 

the source of aqueous sulfate for the more acidic samples. A 

few samples show higher calcium:sulfate ratios than would 

be expected for gypsum/anhydrite dissolution, indicating the 

importance of calcite dissolution. The sulfur isotopes combined with calcium:sulfate ratios prove to be a simple and useful way to classify these waters into three main end members: 

gypsum/anhydrite-dissolution dominated, pyrite-oxidation 

dominated, and calcite-dissolution dominated. Attempts to 

relate the δ18OSO4

 in the pyrite-oxidation dominated waters to 

δ18OH2

O were not successful. Most likely several other processes affect δ18OH2

O without affecting the δ18O SO4

, including 

evaporation and the mixing of different δ18O waters along the 

flow path before they were sampled.

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