Dames & Moore, 1999 - USDA Forest Service

pH Control on Precipitation~Dissolution
Change in pH is probably the most important mechanism controlling the initial fate of metals in mine
.drainage. As described in Section 6.2.1, mine drainage often has a pH of less than 4 due to oxidation of
sulfide minerals. As the acidic leachates contact buffering sources (minerals) and mix with alkaline
waters, the pH increases. The stability of minerals is defined in part by pH. For example, the stability of
aluminum hydroxide in terms of pH can be defmed based on the reaction:
The equilibrium constant for the reaction is:
This defines the relationship between pH (-log an+) and aluminum activity (aA3+) which in turn can be
used to calculate part of the pH-activity diagram shown in Figure 6-3.4. Note that activities (a) are
determined from concentrations adjusted using activity coefficients (y), for example:
Other reactions define other parts of the curve:
41 =Y% (6-20)
AI(OH)2' + H20 o AI(OH)3 + Hf (6-2 1)
.-.'. . . AI(oH)" + 2H20 AI(OH)3 + 2Hf (6-22)
.! .y.- ..,.
The outcome of this diagram is that as an acidic solution evolves in pH by contact with aluminum-
: :: containing minerals, pH increases and aluminum also increases (see hypothetical trajectory). '~ventuall~,
c. the stability line for aluminum hydroxide is reached and aluminum hydroxide precipitates. As long as
aluminum hydroxide is in contact with the solution, the trajectory then follows the curve. down as pH
'increases further (for example, if the water mixes with another alkaline water). Aluminum concentrations
decrease as aluminum hydroxide precipitates. This is a very common effect at mine sites as demonstrated
by the occurrence of white precipitates within the.1500-level main portal drainage (Figure 6.1-la). Figure
6.3-4(a) shows aluminum concentrations on a log scale to illustrate the curve. Figure 6.3-4(b) shows the
&me diagram on an arithmetic scale to illustrate that once the curve is .intersected, and pH continues to
increase, ,aluminum concentrations decrease very rapidly, becoming "undetectable" by pH 5.
similar curves can be drawn for other common precipitates such as red ferric hydroxide (Figure 6.3-5)
and green basic copper carbonate. When carbonates fonn, the partial pressure of carbon dioxide also
affects the pH arid the type of mineral that fonn. For example, under atmospheric conditions, the stable
basic copper carbonate is malachite, but at higher carbon dioxide pore pressures, azurite will form.
The pH changes therefore control the fate of dissolved potential contaminants by causing them to
precipitate. Reversals in pH also cause dissolution by the same processes.
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1769MoS-019Uuly 27,1999;4:11 PM;DRAFT FINAL RI REPORT