Geochemistry of 

Ground Water in Mine 

Drainage Problems 

GEOLOGICAL SURVEY PROFESSIONAL PAPER 473-A 

,,. ,. 

Geochemistry of 

Ground Water in Mine 

Drainage Problems 

By IV AN BARNES and F. E. CLARKE 

MINING HYDROLOGY 

GEOLOGICAL SURVEY PROFESSIONAL PAPER 473-A 

UNITED STATES GOVERNMENT PRINTING OFFICE, WASHINGTON : 1964 
UNITED STATES DEPARTMENT OF THE INTERIOR 

STEWART L. UDALL, Secretary 

GEOLOGICAL SURVEY 

Thomas B. Nolan, Director 

For sale by the Superintendent of Documents, U.S. Government Printing Office 

Washington, D.C. 20402 
CONTENTS 

Abstract------------------------------------------- Introduction ______________________________________ _ 

Role of air and organisms ___________________________ _ 

Oxidation-reduction, a major factor __________________ _ 

Possible anaerobic reactions _______ -------------------

Page 

A1 

1 

1 

2 

3 

Geochemical implications _________ -_-----------------

Chemical effects of neutralization practices ____________ _ 

SummarY------------------------------------------

References-----------------------------------------

TABLES 

Page 

A4 

4 

5 

5 

TABLE 1. Maximum range concentration of Fe+B _________________________ ------------------------------------------ A2 

ILLUSTRATIONS 

FIGURE 1. A diagram from theoretical considerations showing postulated properties of mine water_______________________ A3 

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MINING HYDROLOGY 

GEOCHEMISTRY OF GROUND WATER IN MINE DRAINAGE PROBLEMS 

By I VAN BARNES AND FRANK E. CLARKE 

ABSTRACT 

From information readily available in the literature, the 

compositions of mine waters are considered and possible reactions 

that yield such compositions are examined. Both matter 

balances and thermodynamic considerations indicate that 

reactions taking place in the mines probably do not involve 

significant amounts of dissolved oxygen. The reactions do, 

however, involve water and coexisting geologic materials 

(minerals and probably coal if present), which, in the absence 

of side reactions with carbonates, yield an acidic and strongly 

reduced solution. In the absence of active mine ventilation, 

mine water is an anaerobic solution. Major implications of 

the theoretical model bear on the treatment processes feasible 

for reducing the effect of acid mine water on surface-water 

resourcA: 

INTRODUCTION 

Extensive literature exists on the subject of acid 

coal-mine waters and includes comprehensive reviews 

by Ash and others (1951, 1957), Moulton (1957), 

Hanna and others (1961), and Lorenz (1962). Numerous theories, argun1ents, and contradictions can be 

found on all phases of the subject, including causes, 

mechanisms, treatments, biologic agencies, disposal, 

prevention, and even whether or not such a problem 

exists. 

Coal-mine water contains iron and sulfate in solution, and the water is generally acid in the absence of 

limestone. No obvious method for completely sealing 

mines, especially underground workings, exists. If 

mine workings could have been isolated hydrologically, 

operators, who must pump as much as 23 tons of water 

per ton of coal from the anthracite mines of Pennsylvania (Ash and others, 1957), could have saved a 

great amount of money. 

ROLE OF AIR AND ORGANISMS 

The ultimate source of water in mines is some form of 

atmospheric precipitation. This precipitation is assumed to be saturated with atmospheric oxygen upon 

its arrival at the land surface. A further assumption 

is that no oxygen is lost or consumed by organisms or 

organic material before contacting pyrite (Fe82) in or 

near the coal (Lorenz, 1962; Moulton, 1957). One of 

the postulated reactions for the oxidation of pyrite 

(or marcasite, another form of Fe82) is 

2FeS2+702+2H20=2FeS04+2H2S04. (1) 

The solubility of atmospheric oxygen in water is 15 

parts per million (ppm) at 0° C (Coker, 1954, p. 24). 

The ferrous-iron concentration in the mine water from 

reaction 1 could only be 7.3 ppm. It might be argued 

that if the oxygen were not supplied as a gas dissolved 

in water but as air flowing through the mine openings, 

the oxygen from the air would dissolve in the mine water 

and react with iron disulfide to produce more than 7.3 

ppm ferrous iron in solution. Such an argument is 

not entirely plausible, however, because aqueous 

solutions which contain detectable ferrous iron (0.01 

pp1n by normal analytical methods) oxidize when 

exposed to air and yield ferric iron either in solution 

or in a precipitate. 

Iron oxy-hydroxide ("rust") staining is found where 

mine water overflows onto the land surface. Braley 

(1954, p. 32-33) gives empirical evidence for the complete oxidation of Fe+2 in acid mine effluent upon exposure to air. If air tends to oxidize the ferrous ion in 

solution, the concentrations of ferrous ion (commonly 

a few hundred to a few thousand parts per million) 

reported in mine water (Braley, 1954) probably could 

not be attained by the action of air. Furthermore, in 

the presence of a large quantity of organic material 

(coal), a rapid depletion of oxygen would be expected 

owing to the oxidation of the organic material. Because 

the oxygen concentration in theminewater can be expected 

to be low, another parameter, the redox potential, must 

be introduced for measuring the oxidation state of the 

n1ine water. The redox potential is here defined as the 

potential of a noble-metal electrode in the solution as 

measured against a hydrogen electrode (for which 

E 0 =0), with positive values more oxidizing (European 

convention). Germanov and others (1959) point out 

that significant measurements of dissolved oxygen and 

redox potential are mutually exclusive. If dissolved 

Al 
A2 ~ING HYDROLOGY 

oxygen is detectable, the redox potential is not reproducible-that is, the oxygen electrode is irreversible. 

If the dissolved-oxygen concentration is below the 

threshold of detection, the redox potential may be a 

quantitative measure of the oxidation state of the 

solution. Barnes and Back (1964) show that when 

organic material is in contact with ground water, reliable 

redox potentials can be measured and interpreted in 

terms of the solubility of iron species. 

A number of authors (Lorenz, 1962; Hanna and others 

1961; Moulton, 1957; Braley, 1954; and Leathen, 1952) 

stress the role of organisms in producing acid in mines. 

Life processes require energy, and organisms in systems 

closed to radiant energy must rely on chemical energy 

available in the system for their life processes. In 

utilizing the chemical energy the organisms tend to 

drive the chemical species toward a more stable state. 

Hence, in the absence of an external energy source, 

such as sunlight, the organisms serve as catalysts. A 

catalyst increases the rate of a reaction and frequently 

produces a particular product, but the products are 

always more stable than the reactants. The role of 

organisms in the mine-water environment is to hasten 

the attainment of a more stable state, but, given sufficient time, the same final (equilibrium) state would 

result by either organic or inorganic paths. 

Attempts to apply bactericides in mines have not 

been successful (Lorenz, 1962, p. 21), and, in terms of 

prevention, applying bactericides over the enormous 

area of recharge of the mines probably would not be 

feasible. 

OXIDATION-REDUCTION, A MAJOR FACTOR 

Barnes and Back (1964) show that the maximum 

concentration of ferric ion in ground waters can be 

calculated from the solubility of Fe(OH)3 by using the 

data for 25° 0 cited by Hem and Cropper (1959). The 

reaction is dependent on the pH of the solution, as 

(2) 

for which the equilibrium constant is 

(3) 

where a is the effective or thermodynamic concentration. Defining 

pH= -log aH+t, (4) 

we find 

(5) 

By using Hem's solubility constant and appropriate 

conversion factors, the activity of ferric ion can be 

computed from 

aFe+3(ppm) =3.30X109-3PH. (6) 

The activity or effective concentration of an ion can 

be converted to the concentration as determined by 

chemical analysis through the definitional equation 

a ion 

'Yion 5 C ion 1 (7) 

where 'Yion is the activity coefficient of the ion and aon 

is the concentration of the ion by chemical analysis. 

For ideal solutions, 'Yion = 1. 

Useful activity coefficients of ions can be computed 

for solutions as much as 0.1 molal (Klotz, 1950) in 

total ionic strength; beyond this concentration the 

Debye-Hiickel limiting law does not give reliable 

activity coefficients. Hem (1960) has calculated the 

activity coefficient of the ferric ion to total salt concentrations of 0.1 molal; 'YFe+s is between 1.0 and 0.2 

and decreases as total salt concentration increases. 

The range of concentration of ferric ion has been 

calculated for the limits Of reliability of 'YFe+3 aS a 

function of pH (table 1). 

From the solubility of Fe(OH) 3 a maximum redox 

potential (Eh) for mine water can be computed as a 

function of ferrous-iron concentration and pH. The 

pertinent reaction is 

(8) 

where 

e-1 represents an electron. 

The redox potential is computed from the classical 

N erst equation 

where 

E 0 is the standard potential, 

R is the gas constant, 

(9) 

T is the temperature in degrees Kelvin, 

n is the number of electrons in the reaction, and 

F is the Faraday constant. 

TABLE 1.-Maximum range of concentration of Fe+S 

pH 

2-------------------------- 3 _________________________ _ 

4-------------------------- 5 _________________________ _ 

6--------------------------

Ideal solution 

(ppm) 

3. 3X103 

3.3 

3. 3X10-3 

3. 3X10-6 

3. 3X10-9 

At 25° 0, equation 9 reduces to 

Eh=0.771 +0.059 log aFe+a. 

aFe+2 

0.1 molal solution 

(7Fe+3=0.2) (ppm) 

1. 65X104 

1. 65X10 

1. 65X10-t 

1. 65X10-5 

1. 65X1o-s 

(10) 
GEOCHEMISTRY OF GROUND WATER IN MINE DRAINAGE A3 

The range of Eh for mine waters having activities of 

ferrous ion between 10 ppm and 1,000 ppm in equilibrium with Fe (OH)3 has been calculated and is shown 

in figure 1. Although aFe+2=1,000 ppm, such solutions 

would contain ferrous ion concentrations of about 

several thousand parts per million. The bulk of mine 

water is tentatively assumed to have maximum Eh 

values in the range shown. Eh values for various oxygen 

partial pressures are shown in the same figure. The 

Eh values for the oxygen pressures are obtained from 

the equation used by Garrels (1960, p. 109): 

Eh=l.23+0.01475 log1o Po2 -0.059 pH. (11) 

The Eh for one atmosphere of pure oxygen, the normal 

earth's atmosphere, and the oxygen pressure representing the threshold of determination of oxygen from 

the modified Winkler method are given. The lower 

limit of oxygen from the modified Winkler method is 

0.01 ppm 02, which would give a partial pressure of 

oxygen of 1.2X10-3 a.tm at 25°0. 

1.2 

1.1 

1.0 

0.9 

0.8 

0.7 

0.6 

0.5 (f) 

t-

...J '0.4 0 

> 

~ 0.3 

~ 

w 

0.2 

0.1 

0.0 Fe (0H)3 

-0.1 

Pyrite FeS2 

-0.2 

Fe (0H) 3 -FeS 2 •••••• Fe (0H) 3 -FeS 2 

-0.3 boundary for ·.. / boundary for 

!!so4 = 100 ppm /· ...... <tso4 = 1000 ppm 

-0.4 

-0.5 1 

pH~ 

FIGURE 1.-A diagram from theoretical considerations showing postulated properties 

of mine water. 

Also appearing in figure 1 are two boundaries representing stability between pyrite and ferric hydroxide. 

The higher boundary is for water having a sulfate 

·activity of 1000 ppm, and the lower boundary is for 

. water with a sulfate activity of 100 ppm. Mine water 

usually lies in this range. The reaction along the 

boundary is 

FeS2+11 H20=2 S0;2+Fe(OH) 3+19 H+~+15 e-t, (12) 

for which, at 25°0, by using the data of Latimer (1952, 

p. 12, 39, 221), 

An order of magnitude of change in sulfate activity 

shifts the boundary only 0.009 volt. The boundaries 

are the highest possible in terms of Eh; any less hydrous 

ferric phase would be more stable and would shift 

the boundary down in Eh. Any less stable sulfide 

(marcasite or "amorphous pyrite") would also move 

the boundary downward. The partial pressure of 

hydrogen along the boundary is much higher than the 

partial pressure of oxygen (Garrels, 1960, p. 108). 

Inevitably, as long as the water remains in contact with 

pyrite, the Eh of the water decreases, although at an 

unknown rate, toward the pyrite stability limit. 

Bacteria, if present, merely catalyze the approach to 

equilibrium. 

POSSIBLE ANAEROBIC REACTIONS 

The information above and in figure 1 indicates 

clearly that aqueous solutions in mines may contain 

less than detectable amounts of dissolved oxygen but 

easily detectable amounts of ferrous ion at redox 

potentials far lower than those of oxygen-containing 

solutions. As a consequence, some other reaction 

must be substituted for the one given (equation 1). 

The following equation is suggested: 

FeS2+8 H20=Fe+2+2 S0r2+2 H+~+7 H2• (14) 

The hydrogen may be effectively consumed by reactions such as partial hydrogenation of the nonaromatic 

coal fraction, a suggestion by Irving Breger (oral 

communication, 1963). Solution of the hydrogen in 

the coal should be considered as a possible sink for 

whatever hydrogen is produced. A perhaps less 

attractive suggestion is the possible action of the hydrogen utilized by organisms during the production of 

methane. Such reactions are known in some environments (ZoBell, 194 7), and possibly the same type 

reaction may be found during further study of mine 

water, although coal may be an unfavorable substrate. 

In any event, for pyrite to oxidize and give sulfate 

ions, some other substance must be reduced. Dissolved oxygen alone cannot conceivably be the oxidizing agent, nor is it plausible that the organic matter 

(coal) present is further reduced by the oxidation of the 
A4 MINING HYDROLOGY 

pyrite. The only obvious alternative is to reduce water 

as is suggested in equation 14. The fate of the hydrogen is not known, although it may escape as methane 

or some other organic constituent, or react with oxidized species such as Fe20 3 to produce Fe+2 in solution 

and water. 

The observation of wide pH ranges in mine water 

leads to the obvious conclusion that other reactions 

must occur, such as the reaction on siderite (FeC03) 

which is reportedly found in the coal measures: 

FeCOa+2 H+I=Fe+2+H20+C02. (15) 

Calcite (CaC03) would, of course, react in a manner 

analogous to the reaction of siderite. Clays and other 

silicates may be expected to dissolve as the hydrogenion concentration increases, and concomitant high aluminum concentrations result. As the mine water moves 

into an aerated environment, the ferrous ion oxidizes as 

follows: 

(16) 

requiring a reduction reaction involving other species, 

such as for oxygen, 

(17) 

Combining equations 16 and 17, pH apparently would 

rise, as 

for each ferrous ion oxidized, one hydrogen ion reacts. 

The difficulty is that the solubility of ferric ion is 

strictly limited by pH. From equation 2, 

(19) 

as each ferric ion precipitates, three hydrogen ions are 

produced. 

A strict relation between total iron in solution and 

pH does not exist. The relation depends on whether 

or not a side reaction with a carbonate has taken place, 

how far the oxidation of the initially reduced mine 

water has progressed, and how much ferric ion has 

precipitated. 

GEOCHEMICAL IMPLICATIONS 

There are some significant implications in the view 

of acid mine water presented here. If the suggested 

reactions dissolving pyrite or its polymorphs are 

correct, sealing a mine to prevent ingress of air may not 

solve mine-acid problems satisfactorily. In some 

mines the exclusion of air may lead to the discharge of 

mine water laden with ferrous iron which oxidizes and 

produces the two undersirable effects shown in equation 

20. 

(20) 

First, the precipitation of larger amounts of ferric iron 

would discolor the bed, smother bottom-dwelling 

organisms, and lower the pH of the water. Second, 

the oxidation of ferrous iron tends to decrease the 

amount of oxygen dissolved in the stream and, thereby, 

to decrease the amount of oxygen available to oxygenconsuming organisms. 

Because the reaction that dissolves pyrite causes 

lower redox potentials and higher ferrous and sulfate 

ion concentrations as equilibrium is approached, delaying the passage of water through the mine by partial 

sealing might conceivably aggravate the problem of 

treating the effluent by allowing the solutions to become 

more concentrated. 

As mentioned previously, no practical method of 

hydrologically isolating mines has been found. Exploiting the resource increases the permeability of the 

entire geologic section in underground mining. An 

inevitable consequence of underground mining is 

gradual failure of the overlying rock partly by plastic 

deforn1ation and partly by brittle failure resulting in 

cracks. The failure of old mines in the northern 

anthracite region of Pennsylvania has resulted in 

several feet or more of subsidence. To seal such a 

hydrologically open system that has miles of subsurface 

workings is impossible. As a consequence, ground 

water can flow preferentially through the mines and be 

free to contact pyrite. Thus, mine sealing might lead 

to an eventual steady-state discharge of mine water as 

great as though the mine were unsealed but having 

longer contact time and more opportunity to approach 

equilibriurn with respect to sulfides. 

Observations of mine acid production neither prove 

nor disprove the theoretical geochemistry presented 

here. There are reports (J. E. Kinney, written communication, 1963) of mine-sealing projects during 

which high initial acid production rates declined with 

time as though impoundment inhibited the acidproduction process. On the other hand, the literature 

(Braley, 1954) indicates that certain sealed mines 

yield solutions having much higher acid and iron concentrations than are common for open mines. Such 

inconsistencies indicate the need for better integration 

of theoretical studies, laboratory experiments, and 

field observations. 

CHEMICAL EFFECTS OF NEUTRALIZ!TION PRACTICES 

Chemical neutralization within a mine would not 

significantly alter the theoretical model presented. If, 

for example, limestone were dumped into a mine, the 

reactions would be the same as if carbonate minerals 

were naturally present, as 

(21) 
GEOCHEMISTRY OF GROUND WATER IN MINE DRAINAGE A5 

The reaction would raise the pH, and precipitation of 

ferric hydroxide would be expected. The addition of 

limestone in the mine would, in principle, amount to 

replacing the hydrogen ions and ferrous ions with an 

equivalent amount of calcium ions. The water would 

be very hard, and the hardness, being noncarbonate, 

would be difficult to remove; but the water might not 

become simply a neutral calcium sulfate solution because the neutralization may not stop the pyrite from 

dissolving. If it did not stop the pyrite from dissolving, 

the mine water would still tend toward the phase boundary of Fe (OH)3 and pyrite but would increase in calcium 

and sulfate concentrations. Direct treatment of minewater effluent through a trough containing coarse limestone has been attempted (Lorenz, 1962, p. 10), but this 

process failed because ferric hydroxide precipitated on 

the limestone and armored it against further solution. 

In attempts to decrease the corrosiveness of mine 

water and to improve the quality of water discharged 

into surface-water resources, coal-mine operators have 

used various liming treatments-using slaked lime and 

powdered limestone (Ash and others, 1951, 1957)-for 

neutralization of acid mine waters. Such treatment is 

continued only as long as the mine or breaker is in 

operation and, of course, is useful only where a point 

discharge of mine water is involved. 

For small scattered inflows of mine drainage in to streams 

(seepage), the neutralization methods suitable for point 

sources may not apply. The recovery of acid streams 

to more natural conditions depends on the geology and 

hydrology of the particular drainage basin. It is 

well known that the acid streams recover in the course 

of travel as the iron precipitates and the acid is diluted 

and neutralized by contact with alkaline minerals in the 

streambed and with diluting waters. Unfortunately, 

most streams in anthracite coal fields and in some 

bituminous coal fields are in contact only with sandstone, shale, and coal. Although streams may have 

high pH values, the buffering capacity of the streams 

are low, as bicarbonate concentrations are about 10-50 

ppm. By contrast, bicarbonate concentrations in 

streams in areas underlain by limestone are from 50 to 

200 ppm, and the buffer or neutralization capacity is 

proportionately higher. As a consequence, the recovery 

mechanism of streams in the coal fields is often mostly 

dilution, and neutralization does not occur until tributaries from areas underlain by calcareous rocks enter 

the acid stream. 

An example of natural neutralization occurred where 

Maiden Creek discharges into the Schuylkill River 

above Reading, Pa. (Ash and others, 1951, p. 40). 

While having a discharge that was three times its normal 

flow, the Schuylkill received enough bicarbonate from 

Maiden Creek to raise the pH of the river from 4.4 to 

5.9; the amount of influent bicarbonate was equivalent 

to a maximum of 70 tons of limestone a day, and it 

played a major role in neutralizing the acid accumulated from 880 square miles of watershed (U.S. Geol. 

Survey, 1960) and 45 miles of mainstem channel when 

the flow of the river was three times normal. Thus, 

it may be feasible to use quarry-run limestone for 

raising the buffer capacity of parts of the coal-field 

drainage and beneficiating reaches of streams which 

are now lost for recreational or sewage consumption 

purposes because of an acid condition. In attempting 

such neutralization, placing the limestone upstream 

from the coal-mine effluent or in upstream iron-free 

tributaries should be more effective. If limestone is 

placed in surface water rich in ferrous ions, the precipitation of the ferric hydroxide would tend to insulate 

the limestone and block the desired reaction. 

Neutralization admittedly is not an ideal solution 

to the mine-waste problem. However, less orthodox 

treatment3, such as removal of sulfate ion by an 

electrolytic process using lead electrodes, probably 

cannot be justified in the present situation. 

SUMMARY 

Geochemical considerations of the mine-drainage 

problem indicate that air is not the only major source 

of mine acid; so, mine sealing may not provide a completely 

satisfactory solution to the problem. Indirect neutralization by use of limestone appears to have much 

merit as an interim corrective treatment. Theoretical 

geochemistry should be considered carefully in future 

reBearch on mine sealing or other inhibitive or corrective 

processes because the theoretical model provides a 

basic insight into the causes and effects of acid mine 

drainage. 

REFERENCES 

Ash, S. H., Dierks, H. A., and Miller, P. S., 1957, Mine flood 

prevention and control: U.S. Bur. Mines Bull. 562, 100 p. 

Ash, S. H., Felegy, E. W., Kennedy, D. 0., and Miller, P. S., 

1951, Acid-mine-drainage problems, anthracite region of 

Pennsylvania: U.S. Bur. Mines Bull. 508, 72 p. 

Barnes, Ivan, and Back, William, 1964, Geochemistry of ironrich ground waters of southern Maryland: Jour. Geol. (in 

press). 

Braley, S. A., 1954, Summary report: Mellon Inst. Indus. Research, Pennsylvania Dept. Health, Indus. Fellowships 1-7, 

279 p. 

Coker, Robert E., 1954, Streams, lakes, ponds: Chapel Hill, 

N.C., North Carolina Univ. Press, 327 p. 

Garrels, Robert M., 1960, Mineral equilibria at low temperature 

and pressure: New York, Harper & Bros., 254 p. 

Germanov, A. L., Volkov, G. A., Lisitsin, A. K., and Serebrennikov, V. S., 1959, Investigation of the oxidation-reduction 

potential of ground waters: Geochemistry, no. 3, p. 322-329. 

Hanna, G. P., Jr., and others, 1961, Analysis of fundamentals of 

acid mine drainage: Ohio State Univ. Eng. Expt. Sta. Water 

Resources Center Rept. 175, 76 p. 
A6 ~~G HYDROLOGY 

Hem, J. D., 1960, Restraints on dissolved ferrous iron imposed 

by bicarbonate, redox potential, and pH: U.S. Geol. Survey 

Water-Supply Paper 1459-B, 55 p. 

Hem, J. D., and Cropper, W. H., 1959, Survey of ferrous-ferric 

chemical equilibria and redox potentials: U.S. Geol. Survey 

Water-Supply Paper 1459-A, 31 p. 

Klotz, Irving M., 1950, Chemical thermodynamics: Englewood 

Cliffs, N.J., Prentice-Hall, 369 p. 

Latimer, W. M., 1952, Oxidation potentials: Englewood Cliffs, 

N.J., Prentice-Hall, 392 p. 

Leathen, William W., 1952, Special summary report on microbiological studies of bituminous coal mine drainage: Mellon 

Inst. Indus. Research, Pennsylvania Dept. Health, Indus. 

Fellowship 326B-6, 54 p., and supplement, 9 p. 

0 

Lorenz, Walter C., 1962, Progress in controlling acid mine 

water-a literature review: U.S. Bur. Mines Inf. Circ. 8080, 

40p. 

Moulton, E. Q., ed., 1957, The acid mine-drainage problem in 

Ohio: Ohio State Univ. Eng. Expt. Sta. Bull. 166, 158 p. 

ZoBell, Claude E., 1947, Microbial transformation of molecular 

hydrogen in marine sediments with particular reference to 

petroleum: Am. Assoc. Petroleum Geologists Bull., v. 31, 

no. 10, p. 1709-1751. 

U.S. Geological Survey, 1960, Compilation of records of surface 

waters of the United States through September 1950, Part 

1-B, North Atlantic slope basins, New York to York River: 

U.S. Geological Survey Water-Supply Paper 1302, 679 p. 

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