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Some Chemistry of Acid Mine Drainage (AMD)

by C. Thorsten
CR Scientific LLC
April 2013


The mining of coal deposits and metal ores often results in a phenomenon called "acid mine drainage" (AMD).  This results primarily from the oxidative breakdown of transition metal sulfides, especially pyrite (FeS2).  There is actually a complex series of reactions and side reactions, but the net equation is often written according to Equation 1:

4FeS2 (s)  +  15O2  +  14H2O  ------> 
                              16H+ (aq)  +  8SO42- (aq)  +  4Fe(OH)3 (s)

It is apparent from the above equation that there are two moles of H+ for every mole of SO42-.   Thus, the breakdown of pyrite can in theory provide two moles of sulfuric acid for every mole of FeS2.   However, because ferric hydroxide is soluble in acids, there is an equilibrium between aqueous Fe3+ and solid iron(III) hydroxide (Equation 2).   Notice that in the reverse direction, we actually have the hydrolysis reaction for Fe3+:

Fe(OH)3 (s)  +  3H+ (aq)   <-------->     Fe3+ (aq)  +  3H2O

Dilution of the liquid, or contact with alkaline materials such as carbonate rocks, shifts Equation 2 toward the left, causing precipitation of ferric hydroxide.  As long as the pH of the solution remains low (< 3.5), the equilibrium will be mostly to the right.  Hence, the net reaction (Equation 1) could be expressed in the form of Equation 3:

4FeS2 (s)  +  15O2  +  14H2O  ------> 
                     4H+ (aq)  +  8SO42- (aq)  +  4Fe3+ (aq)  +  12H2O

Again, dilution or contact with carbonate rocks will yield a precipitate of ferric hydroxide.   In reality, the "ferric hydroxide" is a complex mixture of hydrated ferric oxides and oxyhydroxides (e.g., FeOOH), known collectively as "HFO" for Hydrous Ferric Oxides.  A chief constituent of HFO is ferrihydrite;  an approximate formula is 2Fe2O3 H2O, but the amount of water varies.  Ferrihydrite's gross appearance is that of an amorphous, earthy or gelatinous mass. 
AMD-polluted streams may have some iron (II) in solution.  The original pyrite degradation reactions (not shown) initially yield an acidic solution of ferrous sulfate, FeSO4.   This is normally converted mostly to iron (III) by atmospheric oxygen, especially after dilution.  However, organic matter or lack of aeration can favor the presence of Fe2+.  This, in turn, holds more iron in solution up through pH values as high as 8.5.
Although the dissolution of Fe(OH)3 uses up some H+ from Equation 1, the solution from Equation 3 is still very strongly acidic.  Fifty to a hundred years ago, many streams in Appalachia were receiving coal-mine runoff that had a pH of zero or very close to it.   The streams themselves could quite easily attain the pH of vinegar (approximately 2.5) or only slightly higher.
Strongly acidic water has disastrous effects on aquatic ecosystems.   Normally, water that runs over limestone rocks (CaCO3) develops a natural bicarbonate buffer system that resists small changes in pH.  Sufficient acidity will overwhelm this, first protonating the bicarbonate and then driving out the resulting H2CO3 as gaseous CO2.  The result is that further addition of acid causes a precipitous drop in pH, which of course destroys most of the remaining life in the stream.   Compounding the problem is the fact that the hydrous ferric oxides form a layer on which virtually nothing can grow.  This red-orange to ochre-yellow sludge is known as "yellow boy"  (Figure 1):

"Yellow Boy"

Figure 1. A coating of hydrous ferric oxides (HFO) can prevent plants and other aquatic life from growing.  This stream flows through an area last mined over eighty years ago;  the photo was taken in 2013.

The chemistry of AMD is perfectly suited to concentrate its damage on streams and waterways.  First, the very low pH of coal mine runoff (zero to 2.5) holds aqueous Fe(II) and Fe(III) in solution.  Then, when the runoff meets a stream and becomes diluted or contacts carbonate rocks, the slightly higher pH causes Fe(OH)3 to precipitate.  Even if the acid runoff should cease, this layer makes it difficult for an ecosystem to flourish again.
An additional problem with AMD is that it often contains aluminum and lesser amounts of other pollutants, such as arsenic.  Clays and other minerals leach some aluminum under strongly acidic conditions, since the sulfuric acid from AMD readily forms aluminum sulfate.  Furthermore, aluminum forms soluble sulfate complexes (e.g., AlSO4+).  Like iron and other transition metal ions, Al3+ undergoes hydrolysis in water to yield an acidic solution.  That is because aluminum sulfate, like iron sulfate, is the salt of a weak base and a strong acid. 
As in the case of Fe3+, an increase in pH can cause aluminum to precipitate in the form of hydroxides and hydrous oxides.   Very alkaline water can actually re-mobilize aluminum, however, since the hydroxide is amphoteric.  Aluminum precipitates most effectively at pH 6.5 to 7.  Nevertheless, such near-neutral waters can still contain considerable amounts of both sulfate and iron.  AMD therefore presents remediation challenges that go beyond simply neutralizing the water.
AMD-polluted streams often look like something we'd expect on the surface of a distant moon.  There is not a plant, not a fish, not anything familiar.  Most kinds of life cannot flourish in such waters.    Even after pH neutralization, there are lasting effects.   This is the legacy of acid mine drainage.


Figure 2: We already know the reddish-brown coloration is due to iron (III).  Note, however, the bluish-green color in the deeper portions of the stream.  Something interesting is happening here. 

We will now look at a few informal experiments and test methods, together with some further comments on the chemistry of AMD.

pH Tests 

Some pH measurements may help us explain at least some of what we observe at the site.  That is because, to a large extent, the pH determines what ionic species, and how much of them, the water will have.  This fact, by the way, is central to the design of passive AMD treatment systems.
We can connect the measured pH to other observed facts, as well;  these may yield some clues about the history of an AMD-polluted stream.  Is the water flowing over carbonate rocks?  Has any remediation been performed in the area?  How does the stream pH compare to other branches or creeks that are not polluted?  Do the rocks have a yellowish-orange (iron) or whitish (aluminum) coating?  Is there a place in the stream above which the pH changes noticeably?
A creek with "yellow boy" but which has near-neutral pH (6.5 to 7) is likely to have mine sites that have already been treated with alkaline material, or perhaps the creek has been flowing over limestone for some distance.  The water will probably contain very little soluble aluminum, but it may still contain unacceptable amounts of iron.  At neutral pH, we can expect the water to have developed a functioning bicarbonate buffer system, since it can be assumed that carbonate rock was used to neutralize the AMD in most situations.
A creek with much more acidic pH (2.5 to 4) will lack a functioning bicarbonate buffer system.  Such a low pH probably indicates an ongoing, nearby source of sulfide mineral degradation.  Note the rock types in and near such a stream. 
If you can find any shallow, slow-moving regions with heavy deposits of Fe(OH)3, check the pH there.  Is the water acidic?  Why or why not?  Is this consistent with Equation 2?


The concentration of sulfate can remain undesirably high even after an AMD-polluted stream has been neutralized with limestone.   However, we might expect the sulfate concentration to be highest in active AMD-producing sites, since there has not been sequestration of sulfur species in the form of either sulfides or poorly-soluble sulfates (e.g., calcium sulfate or basic aluminum sulfate).
Most any chemist knows the classic method to test for sulfate;  it depends on the very low solubility of BaSO4.  Given a water sample with an expected sulfate concentration of, say, 500 ppm, and an electronic balance having readability of 0.001 g and capacity of 30 grams, it is fairly straightforward to determine the minimum amount of sample water that must be collected, as well as the minimum amount of BaCl2 to add.


Ferrous ion oxidizes to ferric in the presence of dissolved oxygen (Equation 4):

4Fe2+ (aq) + 4H+ (aq) + O2 (aq)  ------>  4Fe3+ (aq) + 2H2O

Looking at the net equation (above), we would assume that H+ drives the formation of Fe3+ when O2 is present.  However, this writer knows from experience that in the laboratory,  a pH of 4 or lower will stabilize Fe2+ and will actually favor the reduction of Fe3+, even in the presence of atmospheric O2.  The result is a solution that is practically all ferrous sulfate.  In other words, H+ actually favors the opposite of what is shown in the above equation.  In fact, it is well-known in the literature that acidic solutions stabilize Fe2+;  see for example Stumm and Morgan (1995).
Oxidation of iron (II) does occur in natural environments at fairly low pH. This is not driven by H+ per se.  In fact, while oxidation at low pH is primarily the result of phototrophic bacteria (Hegler et al. 2008), even these organisms do not begin to oxidize Fe(II) until the pH exceeds 5.   Thus, we must treat Equation 4 as a "net reaction" that really speaks of multiple processes.  Again, Fe(II) is actually stabilized by high concentrations of H+,  and therefore we cannot regard H+ as if it drives the reaction. 
Fe3+ in aqueous solution precipitates easily as Fe(OH)3, especially at increased pH.  In streams affected by AMD, precipitation occurs at pH 3.5 or above, rather than 4 or above as we'd encounter in a laboratory.  This is probably due to the increased oxygenation caused by turbulence.   Since "ferric hydroxide" is actually a mixture of hydrous ferric oxides, it stands to reason that oxygen will promote the formation not only of Fe3+ from Fe2+ (Equation 4), but also hydrous ferric oxides from Fe3+.  In practice this is exactly what happens.  Leave a neutral solution of ferrous sulfate in contact with air, and it doesn't simply form ferric sulfate;  it forms hydrous ferric oxides.  Aerate that solution, and a larger amount will form.
Qualitatively, a very high concentration of iron is easy enough to discern in a stream:  simply look for the "yellow boy" coating on the rocks.   However, it is useful to know just how much Fe(III) and perhaps Fe(II) is present.  Water that has flowed with turbulence over rocks and undergone continuous air bubbling should have more Fe(III), whereas deeper water that has been undisturbed, and which may have organic matter, will have significant amounts of Fe(II). The half-reaction

Fe3+ (aq)  +  e-  ------>  Fe2+  (aq)

(Equation 5) in the presence of H2SO4 has a highly positive standard reduction potential (about 0.778 volt).  It will therefore proceed spontaneously if a suitable electron donor is available.  If not oxidized by air, Fe(II) can then be held in solution up until the pH reaches 8.5, at which point greenish Fe(OH)2 will precipitate. 
If the water is at pH 6.5 to 7, and there hasn't been much aeration, it would be a good idea to test for Fe(II), selecting a method that would function in the presence of Fe(III).  There are several tests available, including traditional methods of titration.  For example, total iron can be determined by titration with K2Cr2O7 in the presence of acidified redox indicator (e.g., diphenylamine).
For humans, the US EPA-recommended upper limit for iron in natural waters is 300 mg per cubic meter, which is 0.3 milligrams per liter (i.e., 0.3 ppm).  With AMD, the iron concentration will be significantly higher than 0.3 ppm.  A polluted stream might have a total iron concentration between 10 and 100 ppm, with most of this being Fe(III).  This would influence to some extent the analytical method chosen, although Fe(III) can be reduced to Fe(II) if needed.  Most of the wet chemical methods will function very well at iron concentrations above 1 ppm, though a few have a threshold of 10 ppm or higher. 

Figure 3.  Hydrous ferric oxides may gradually take on a darker, rust-red appearance, especially in areas where the rocks occasionally dry in the sun. The loss of H2O from these hydrous oxides yields α-Fe2O3, the familiar red mineral known as hematite.


The chemistry of iron and aluminum sulfates,  hydroxides, and hydrous oxides is central to the understanding of AMD remediation.   Because modern civilization depends on mining, and because there are so many old mine sites still in need of AMD remediation, these subjects will continue to be important for many generations.  In fact, AMD can occur when there is road construction or other large-scale disruption of sulfide-bearing deposits. 
The chemical reactions of AMD can actually be very complex because of the interplay of so many variables.  For example, even if we know how much of a particular neutralizing agent is added, the future state of a stream can be very difficult to predict.  Each site requires ongoing data-gathering and the assessment of locally-unique conditions such as topography, soil permeability, and so on.  However, all these factors ultimately meet one another in the chemical reactions that take place in the streams and waterways.


Hegler, F., Posth, N., Jiang, J., and Kappler, A.  "Physiology of Phototrophic Iron(II)-Oxidizing Bacteria:  Implications for Modern and Ancient Environments".  FEMS Microbiology Ecology 66(2): 250-260 (November 2008).

Stumm, W and Morgan, J.J.  Aquatic Chemistry: Chemical Equilibria and Rates in Natural Waters. New York:  Wiley-Interscience, 1995.

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