1249 

### COMPARISON OF THREE METHOD S TO MEASURE ACIDITY OF
COAL-MINE DRAINAGE 1

Brent Means 2 and Tiff Hilton
Abstract. Although the Standard Methods 2310 hot peroxide acidity procedure is widely used for measuring the acidity of mine draina ge, little work has been done to determine if “hot acidity” data actually describe the base requirement for neutralization of mine drainage. This study compared three met hods for estimating the acidity of net-acidic waters emanating from the Manor, Millers tein, Ike, and Morris coal mines in Pennsylvania: Standard Methods 2310 hot aci dity titration to pH 8.2 endpoint, cold acidity titration to treatment endpoint pH as high as 11.0, and calculated acidity. The results showed poor agreement between hot acidi ty and calculated acidity for three of the four waters. For two of the waters, Mg hydrolysis during the hot-acidity titration indicated greater acidity than that computed based on pH and dissolved Fe, Mn, and Al. The poor agreement for the other water resu lted from incomplete hydrolysis of Mn during the hot acidity titration. The agre ement between the acidity measured by the treatment acidity titration and the other tw o acidity methods was within 16 mg/L (as CaCO 3 ) for the Manor and Millerstein waters , but greater than 200 mg/L (as CaCO 3 ) for the Ike and Morris waters. The fair agre ement between all methods for Manor and Millerstein is a result of pH, Al, Fe, and Mn being the main source of acidity in the waters. The poor agreement between the acidity methods for the Ike and Morris waters is a result of the treatment acidity titration measuring a large amount of additional “acidity” at high pH from the hydrolysis of Mg and other constituents. While the exact sources of acidity measured by a treatment titration is unknown, results from PHREEQC aqueous speciation calculations showed that the fo rmation of cation-hydroxl complexes in the Morris water at pH 11.0 contributed 40 mg/L of acidity. The authors hypothesize that Mg hydrolysis and the formation of b ase-consuming complexes are the reason why acidity measured by treatment titrations at high pH is often greater than that measured by hot acidity titrations to pH 8.2 or 8.3. The authors also hypothesize that the neutralization of carbonic acid is the reason why the acidity measured by “cold” titrations at low to mid pH is often greater than that measured by hot acidity titrations. The results of this study have practical importance because they show hot acidity titrations should not be used to universally describe the acidity of mine draina ge. This is especially true when estimating the acidity produced when Mg-rich mine drainage is chemically treated to high pH. This study also showed that over treating Mg-rich mine drainage not only increases chemical costs but also increases sludge production.

_________________________________ 1 Paper was presented at the 2004 National Meeting of the American Society of Mining and Reclamation and the 25 th West Virginia Surface Mine Drainage Task Force, April 19-24, 2004. Published by ASMR, 3134 Montavesta Rd., Lexington, KY 40502. 2 Brent Means, Hydrogeologist, U.S. Office of Surface Mining, 415 Market Street, Harrisburg, PA 17011 ( bmeans@osmre.gov , Tel. 717-782-4036); Tiff Hilt on, Mining Engineer, WOPEC Inc., RR2 Box 294B, Lewisburg, WV 24901 ( wopec@charter.net , Tel. 304-645-7633).1250 Introduction

It has long been recognized th at acidity is an im portant measure for characterizing mine drainage (Rose and Cravotta, 1998; ADTI, 2000; Ki rby, 2002). Various methods can be useful to determine the acidity of water and these methods fit into one of three categories: cold acidity, hot acidity, or calculated acid ity. Both cold and hot acidit y (A.P.H.A, 1998) are laboratory procedures that measure acidity by titrating wa ter with a strong base (e.g. NaOH), while the calculated acidity method (Hedin et al ., 1994) determines acidity by using a pH measurement
and dissolved metal concentrations. Coal-min ing professionals in Appalachia have long
considered a Standard Methods 2310 hot acidity titrati on performed to pH 8.3 (American Public Health Association, 1998) as the industry standard for measuring the acidity of mine drainage. This may have developed as a resu lt of regulatory agencies, such as Pennsylvania Department of Environmental Protection (PADEP), requiring that acidity measurem ents (to 8.3) be included as part of the hydrologic data submitted in a mining permit (Pennsylvania Department of Environmental Protection, 2004). Results of hot acidity titrations are used to quantify the severity of mine drainage (Ros e and Cravotta, 1998) and to calcu late limestone consumption in passive treatment systems (Hedin et al ., 1994). Recent work by Kirby and Cravotta (2004)
identifies the sources of acidity measured by a hot acidity titration. Howeve r, little work has been
done to describe the acid ity of mine drainage at titration en dpoints higher than pH 8.3, which is consistent with active treatment methods, or to evaluate the appropriateness of ubiquitously using hot acidity results to describe the acidity of mine drainage. This paper addresses these points by comparing the results of applying cold, hot, and calculated acidity methods to four netacidic discharges emanating from bituminous coal mines in Pennsylvania. First, a background section reviews both the geochemistry of acidity and the th ree acidity methods us ed in this study. Second, a site selection section discusses the reasoning behind choos ing the four mine waters. Third, a discussion section presents the results of the applying the acidity methods to the four mine waters. Finally, practical implica tions of the results are discussed.1251 Background

Overview of Acidity in Mine Drainage Mine drainage commonly contains elevated am ounts of dissolved metals (Rose and Cravotta, 1998). In fact, U.S. Office of Surface Mining’s (OSM) mine drainage inventory shows that 360 of 1470 post-primacy discharges in the Appal achian coal fields have aluminum, iron, or manganese concentration of greater than 50 mg /L (unpublished data). At favorable pH conditions, Al 3+ , Fe 2+ , Fe 3+ , and Mn 2+ metals participate in hydrolysis reactions and produce acidity:

+ + + = + H OH Al H Al 3 ) ( 0 3 3 2 3 (1)

+ + + = + + H OH Fe O H O Fe
s 2 ) ( 5

. 2 25

. 0 ) ( 3 2 2 2 (2)

+ + + = + H OH Fe H Fe 3 ) ( 0 3 3 2 3 (3)

+ + + = + + H MnO O H O Mn
s 2 5

. 0 ) ( 2 2 2 2 (4)

Fe 2+ and Mn 2+ are redox dependent and are shown in e quations (2) and (4) as precipitating as chemically oxidized solids, which are typically fo rmed when passively treating mine drainage or after prolonged exposure of mine waters to atmospheric conditions in surface-water environments. During chemical treatment of mine drainage, Fe 2+ and Mn 2+ can precipitate as chemically reduced solids (Equations (5) and (6)); however, the acidi ty generated is identical to that expressed in Equations (3) and (4).

+ + + = + H OH Fe H Fe 2 ) ( 0 2 2 2 2 (5)

+ + + = + H OH Mn H Mn 2 ) ( 0 2 2 2 2 (6)

In addition to elevated metals, many Appalach ian discharges are characterized by low pH. OSM’s mine drainage inventory shows that 430 of 1470 post-primacy discharges in the1252 Appalachian coal fields have pH < 3.5 (unpublished data). Acidity from low pH (e.g. pH<3.0) is
substantial and can be ca lculated by Equation (7): 

## ( ) pH LofCaCO mg pHacidity
− × = 10 000 , 50 / 3 (7)

In addition to metal acidity and pH acidity, bisulfate (HSO 4 - ), carbonic acid (H 2 CO 3 ), and hydrogen sulfide (H 2 S) are known to contribute to the acidity of fresh mine drainage (Jageman et al ., 1988; Rose and Cravotta, 1998). Payne a nd Yeates (1970) also documented positive interferences from magnesium to hot acidity values.

Acidity Methods

Calculated Acidity Equations (1) through (6) show the amount of acidity produced through hydrolysis is proportional to the concentrations of dissolved metals expresse d as mole equivalents. For instance, Equation (1) s hows that 3 moles of acidity are produced for every 1 mole of Al 3+ participating in a hydrolysis reaction. This stoichiometry permits the development of an equation that uses dissolved aluminum concen tration to estimate acidity generation from hydrolysis and precipitation.

        × × = + + 27 3 50 3 3 Al

acidity Al (8)

where Al 3+ Acidity is the amount of acidity in mg/L of CaCO 3 , Al 3+ is dissolved Aluminum concentration in mg/L, and the values of 27 and 50 are the mol ecular weight of aluminum and the conversion factor for expressing acidity as CaCO 3 , respectively. Hedin et al., (1994) combined Equation (7) with Equation (8) to yield an equation capable of estimating both metal (Fe, Al, and Mn) and pH acidity.

     

× +         × +         × +         × +         × × = − + + + + pH Mn Al Fe Fe Acidity 10 1000 55 2 27 3 56 3 56 2 50 2 3 3 2 (9)1253 where Acidity = metal (Fe, Al, and Mn ) and pH acidity as mg/L of CaCO 3 and Fe 2+ , Fe 3+ , Al 3+

, and Mn 2+ are the analytical concentrations of dissolved metals in mg/L. Equation (9) uses result s from a field-measured pH coupled with results from a laboratory analysis for metals (dissolved) from a field-filtered sample. The accuracy of the acidity estimation for low -pH discharges is highly sens itive to the accuracy of the pH reading. For example, the pH contribution to acidity for pH 3.0 and pH 2.7 water is 50 mg/L (as CaCO 3 ) and 100 mg/L (as CaCO 3 ), respectively. This example illustra tes the importance of using a calibrated pH meter to measure field pH when using Equation (9). A la boratory-measured pH should not be substituted for a field-measured pH since a lab pH could be unreprese ntative if pH-altering chemical reactions, such as hydrolysis and exsolution of CO 2 , occurred during sample transport (Cravotta and Kirby, 2004). It is also important that metal concentrations used in Equation (9) are from analysis of a field-filtered sample as opposed to an unfiltered or laboratory-filtered sample. It is common practice to consider chem ical species contained in a field-filtered (0.45

μ m) sample as dissolved, while an unfiltered sa mple contains both dissolved and suspended species. The use of metal concentrations from an analysis of an unfiltered sample could result in an artificially high calculated acidity, if the sample contained suspended matter (e.g. metal hydroxides and oxides) that is not actually a source of acidity (already hydrolyzed) but could not be distinguished from dissolved concentrations after sample pr eservation with acid. Table 1 illustrates this point.

Table 1. The effect of using metal concentrations from unfiltered samples in Equation (9)

Results of Lab analysis
for Al 3+ (mg/L) Calculated Acidity from Al hydrolysis (mg/L CaCO 3 ) Sample Filtered 5 28 Sample NOT Filtered 10 56

For the purposes of compari ng the results of th e three acidity methods, net acidity was calculated by subtracting alkalinity from calculated acidity (Equation 9):

Net Calculated Acidity (mg/L CaCO 3 ) = Calculated Acidity – Alkalinity (10)1254 Hot Acidity The adjective “hot,” in hot acidity titration, refers to the boiling of the sample before the water is titrated with a strong base to a speci fied pH endpoint. Several published hot acidity procedures exist (APHA, 1998; ASTM, 1984; U. S.EPA, 1979), however, mo st laboratories in Pennsylvania use either Standard Methods 2310 (APHA, 1998) or EPA 305.1 (U.S. EPA, 1979) procedures (Means, unpublished data). Althoug h the EPA (1979) procedure specifies the hot acidity titration endpoint pH of 8.2, and the APHA (1998) and ASTM (1984) procedures specify the endpoint pH of 8.3, all thr ee methods employ consistent steps. Kirby and Cravotta (2004) offer a thorough discussion of the hot acidity tit ration procedure, resulting in an abbreviated discussion for this paper. The general hot acidity steps are to (1) titrate th e water using a sulfuric acid to pH less than 4.5 to neut ralize alkalinity; (2) add five drops of a 30% solution of hydrogen peroxide (H 2 O 2 ) for metal oxidation; (3) boil the sample for two to five minutes to remove all carbonate species, accelerate oxidati on and precipitation of metals, and then allow the sample to cool to room temperature; (4) titrate the water us ing a strong base to pH 8.3 to neutralize acidity; and (5) calculate hot acidity by s ubtracting the equivalent of acid added from the equivalent of base added and express hot acidity as mg/L of CaCO 3 . For the purpose of this paper, there are two important points in this proc edure. First, the procedure intentionally does not measure carbonate acidity (H 2 CO 3 ). Secondly, step (5) yi elds a hot acidity result that characterizes water as net acidic (positive value) or net alkaline (negative value).

Treatment Acidity The authors developed the treatment acidity procedure as a means to measure the base consumption for actual treatment of raw mine wa ter to effluent limits. A treatment acidity
titration is a variant of a “conventional” cold acid ity titration. There ar e two general differences
between a treatment titration and a co ld acidity titration. First, unlike a conve ntional cold acidity titration, a treatment acidity titration is a group of individual titrations with each titration using a different pH endpoint. The pers on performing the titration determ ines the number of titrations and the range in pH endpoints by studying the raw water chemistry and estimating the approximate range in pH which the water will meet effluent standards (Table 2). At the end of each titration, acidity is calculated and the water is filtered and analyzed for dissolved metal concentrations. This process describes the behavior of metals as the pH of the water is increased1255 and provides the pH where efflue nt standards are met. The s econd difference is the point at which the titration ends. In most cases, a “conven tional” cold acidity titra tion requires chemical be added until pH no longer drifts from the desired endpoint, whereas, a treatment titration is finished immediately after the pH endpoint is r eached and pH is free to drift downwards as hydrolysis reactions continue to occur. This di stinction stems from the difference in goals for each procedure. The goal of a cold acidity titration is to calculat e the acidity of the water at a predetermined pH endpoint, while the goal of a treatment titration is to estimate the acidity
neutralized when alkali chemi cal is added to the water by a full-scale treatment system. Treatment systems add sufficient treatment chemical to meet effluent standards, not a “stable” pH endpoint. As a result, the acidity meas ured from both methodol ogies will differ.

Table 2. These pH ranges do not nece ssarily correspond to the minimum solubility for each species, but rather correspond to the approxim ate ranges in pH needed to meet effluent standards for a specific meta l (based on authors’ experience).

Metal pH range needed to meet metal effluent
limits
Fe 3+ 3.5 – 4.5
Al 3+ 4.5-6.0
Fe 2+ 7.0 – 9.0
Mn 2+ 9.0 – 10.5

The basic procedure for a treatment titration is as follows: (1) P our a liter of raw water into a 2000 ml beaker set atop of a magnetic stirrer. (2) Use a Hach digital titrator and a 5N solution of laboratory grade sodium hydroxide (NaOH) to titr ate the water to the target pH endpoint. Record the amount of NaOH consumed. (3) Filte r approximately 125 ml of the titrated water using a hand pump and a 0.45- μ m filter. Add HNO 3 until the pH of the samp le is less than 2.0 to preserve the sample for laborator y metal analysis. (4) Repeat St eps 1-3 to various pH endpoints until you believe you have captured the point the water met effluent standards for metals concentration. (5) Analyze the acidified samples for metals of interest, such as iron, aluminum, manganese, magnesium, and calcium. (6) Use the labo ratory analysis of metals concentration to determine the point at which the water meets efflue nt limits and calculate the treatment acidity.1256 Water Quality Collection and Analysis

Calculated Acidity The following procedure was used to determin e calculated acidities. In the field, mine drainage was passed through a 0.45- μ m filter using a hand-held vacuum pump. Half of the water was preserved with reagent-grade HCl to pH < 2. 0 for determining ferrous iron concentrations.
The remainder of the water was preserved with reagent-grade HNO 3 to pH < 2.0. Cardan
Laboratories of Northern Cambri a, Pennsylvania determined meta l concentrations using atomic absorption (APHA, 1998). At the time of sampling, pH was measured in the field using a Hydac pH meter calibrated in pH 4.0 a nd 7.0 buffer solutions. A sample was collected, placed on ice, and analyzed for alkalinity within 24 hours us ing a Hach alkalinity titration kit.

Hot Acidity Unfiltered water samples were collected and pl aced on ice until hot ac idity titrations were performed. During the manual hot acidity titrat ions, Cardan Laboratorie s performed additional steps than required by Standard Methods 2310 procedures (APHA, 1998). For example, after the hot peroxide boil a small portion of water was filtered and analyzed for metal concentrations to study the effect of the hot peroxide boil. Also , the water was titrated to pH endpoints 8.2 and 11.0 for the purpose of investigating sources of acidity at higher pH. At each pH endpoint, a portion of water was filtered and analyzed for metal concentrations to identify the sources of acidity. This paper refers to th e hot acidity titration performed to pH 8.2 endpoint as “hot acidity (pH 8.2)” and the hot acidity titration performed to pH 11.0 endpoint as “hot acidity (pH 11.0).”

Treatment Acidity Unfiltered water samples were collected and refrig erated and treatment aci dity titrations were performed within 24 hours. Pennsylvania Bureau of La boratories determined metal
concentrations from the filtered water obtained at each of the different titration endpoints using inductively coupled plasma emissi on mass spectrometry (APHA, 1998).1257 Site Selection The three acidity methods were applied to four net-acidic discharges emanating from Pennsylvania coal mines. The four discharges were selected because they possess a wide-range of water quality, which is desira ble to study different sources of acidity (Tables 3 and 4). The discharges Ike, Millerstein, and Morris are fr om reclaimed surface mines and Manor emanates from an underground mine. Ike was selected becau se of its moderate me tal concentrations and pH. Manor was selected because of its high ferrous iron concentration and low magnesium
concentration. Millerstein was selected because of its moderate pH and low concentration of aluminum, iron, and manganese and moderate con centration of magnesium. Morris was selected because of its high concentrations of aluminum , manganese, magnesium, zinc, and sulfate, and
low concentration of iron.

Table 3. Major cation and anion raw water quality data analyzed by Pennsylvania’s Bureau of Laboratories [concentrations (mg/L) are from a filtered sample; pH is field measured]

Site pH Al Ca Fe 3+ Fe 2+ K Mg Mn Na SO 4 2-

Ike 4.18 5.95 150.0 <0.02 24.0 4.5 134.0 24.0 7.5 1341.2

Manor 2.98 33.0 190.0 0.03 326.0 2.8 57.0 4.7 3.81 1726.2

Millerstein 5.98 0.25 101.0 <0.02 4.12 2.8 113.0 5.2 46.0 1380.0

Morris 3.01 61.0 186.0 <0.02 0.75 3.8 460.0 57.0 18.0 4223.5

Table 4. Trace metal raw water quality da ta analyzed by Pennsylvania’s Bureau of Laboratories [concentrations (mg/L) are from a filtered sample; pH is field measured]

Site pH As Cd Co Cu Ni Se Zn

Ike 4.18 <0.004 <0.01 0.34 0.01 0.48 <0.007 1.032

Manor 2.98 <0.004 <0.01 0.18 0.03 0.294 <0.007 1.285

Millerstein 5.98 <0.004 <0.01 <0.05 <0.01 0.11 <0.007 0.105

Morris 3.01 <0.004 <0.01 1.330 0.033 2.1 0.0074 4.1221258 Results

Table 5 shows the results of Cardan Laboratories analysis of the filtered raw water. These concentrations were used as input into Equation (9) to determine calculated acidity (Table 6).
Tables 7 and 8 show how dissolved concentrations change at vari ous steps during the hot acidity titration. The first column in Table 7 presents the dissolved concentrations for the raw water of Ike discharge, and the second colu mn presents the dissolved concen trations after the hot peroxide
boil. The third and fourth columns show the disso lved concentrations after the completion of the hot acidity titration to pH 8.2 and pH 11.0 endpoints. Tables 9 – 12 show the results of the treatment acidity titrations. These tables show the changes in dissolved concentrations as the raw water is incrementally titrated to numer ous pH endpoints, filtered, and analyzed for
dissolved metals. These tables also show the corresponding acidity for each incremental
titration. The bolded rows in thes e tables represent the titration that resulted in (approximately) meeting effluent standards defined by Pennsylva nia regulations. Table 13 provides a summary of the results for each acidity method.

Discussion of Results

Table 13 shows that the hot ac idity titrations performed to th e endpoint pH of 11.0 produced the highest acidity values, followed by the treatmen t acidity titrations. The table also shows that calculated net acidities and hot acidities (pH 8.2) are more similar than treatment acidities and hot acidities (pH 8.2). The acidity methods that produced similar results are measuring the same sources of acidity, whereas met hods that produced different re sults are measuring different sources of acidity. In order to correctly apply ea ch acidity method in the fu ture, it is important to identify the sources of acidity that each method measures. Equation (9) only accounts for acidity associated with pH, Al, Fe, and Mn components of a sample. Also, this method assumes complete hydrol ysis of Al, Fe, and Mn . Table 13 shows that
calculated net acidities and hot aci dities (pH 8.2) for Ike are vi rtually identical, which suggests both methods measure the same sources of acidity.1259 Table 5. Raw water analyses pe rformed by Cardan Laboratories.

Ike
Dissolved conc. of Raw water (mg/L) Manor
Dissolved conc. of raw water (mg/L)) Millerstein
Dissolved conc. of raw water (mg/L) Morris
Dissolved conc. of Raw water (mg/L) Field pH 4.22 3.02 5.98 3.01
Al 6.73 37.0 1.0 71.6
Ca -- 190.0 -- --
Fe 21.8 326.0 3.22 0.73
Mn 23.6 4.26 5.16 55.6 Mg 134.0 57.0 -- --
SO 4 2- 1341.2 1726.2 1380.0 4223.5

Table 6. Results of using Equation 9 to calculate net acidity.

Discharge

Name Calculated Acidity (mg/L CaCO 3 ) Field Alkalinity (mg/L CaCO 3 ) Net Acidity (mg/L CaCO 3 )

Ike 122.4 2.0 120.4
Manor 845.0 0.0 845.0
Millerstein 20.7 16.0 4.7
Morris 549.4 0.0 549.41260 Table 7. Hot Acidity titration for the Ike and Ma nor discharges [Dashes indicate not determined or not applicable].

Ike Manor

Dissolved
conc. from

raw water

(mg/L) Dissolved
conc. after
H2O2 boil

(mg/L) Dissolved
conc. after Hot Acidity

Titration to pH 8.2 (mg/L) Dissolved
conc. after Hot Acidity

Titration to pH 11.0

(mg/L) Dissolved
conc. from

raw water

(mg/L) Dissolved
conc. after
H 2 O 2 boil

(mg/L) Dissolved
conc. after Hot Acidity

Titration to pH 8.2 (mg/L) Dissolved
conc. after Hot Acidity

Titration to pH 11.0

(mg/L) Lab pH 4.53 3.13 8.2 11.0 2.98 2.43 8.2 11.0
Al 6.73 4.26 0.16 0.05 37.0 30.2 0.26 3.65
Ca -- 146.0 136.0 109.3 190.0 183.4 180.4 112.8
Fe 21.8 0.8 0.01 0.01 326.0 66.3 0.25 0.03
Mn 23.6 23.55 4.03 0.18 4.26 4.1 0.55 0.01 Mg 134.0 133.1 120.4 22.0 57.0 48.7 48.5 0.77
SO 4 2- 1341.2 1237.8 1165.0 1110.4 1726.2 1700.0 1587.5 1468.7 Hot Acidity * -- -- 123.0 720.0 -- -- 880.0 1360.0 Lab Alkalinity* 2.0 -- -- -- 0.0 -- -- --

*Expressed as mg/L of CaCO 31261 Table 8. Hot Acidity titration for the Millerstei n and Morris discharges [Dashes indicate not determined or not applicable].

Millerstein Morris

Dissolved
conc. from

raw water

(mg/L) Dissolved
conc. after
H2O2 boil

(mg/L) Dissolved
conc. after Hot Acidity

Titration to pH 8.2 (mg/L) Dissolved
conc. after Hot Acidity

Titration to pH 11.0

(mg/L) Dissolved
conc. from

raw water

(mg/L) Dissolved
conc. after
H2O2 boil

(mg/L) Dissolved
conc. after Hot Acidity

Titration to pH 8.2 (mg/L) Dissolved
conc. after Hot Acidity

Titration to pH 11.0

(mg/L) Lab pH 6.2 7.44 8.2 11.0 3.01 2.93 8.2 11.0
Al 1.0 0.24 0.11 0.07 71.6 54.71 1.1 0.15
Ca -- 111.2 106.0 99.1 -- 195.4 189.8 109.76
Fe 3.22 0.01 0.01 0.01 0.73 0.64 0.03 0.03
Mn 5.16 3.2 3.2 0.12 55.6 43.0 21.2 0.11 Mg -- 119.8 116.7 39.5 -- 477.4 424.3 9.34
SO42- 1380 1146.8 1098.3 1020.5 4223.5 3523.7 3175.0 2950.0 Hot Acidity* -- -- 16.0 1340.0 -- -- 510.0 2840.0 Lab Alkalinity* 16.0 -- -- -- 0.0 -- -- -- *Expressed as mg/L of CaCO 31262Table 9. Results of treatment aciditytitration for Ike. Dissolved metal concentrations as a function of pH.TitrationpHendpointAl As Ca Cd Co CuFe3+Fe2+K Mg Mn Na Ni Se ZnTreatmentNetAciditymg/LCaCO3AnnualCost totreat aflow of 1gpm**4.18 (raw)5.95 <0.004 150.0<0.01 0.34 0.01 <0.0224.04.5134.00 24.0 7.50.48 <0.0071.0325.51.0 <0.004150.0<0.010.314 <0.01<0.0219.0 4.5 138.0 24.0 25.0 0.485<0.0071.068 39 $416.50.2 <0.004150.0<0.010.289 <0.01<0.0213.0 5.3 138.0 24.0 40.0 0.454<0.0070.936 70 $747.0<0.2 <0.004144.0<0.010.242 <0.01<0.021.6 4.9 132.0 23.0 95.0 0.343<0.0070.486 207 $2188.0<0.2 <0.004141.0<0.010.177 <0.01<0.020.42 4.4 130.0 21.0 113.0 0.281<0.0070.435 232 $2449.0<0.2 <0.004141.0<0.010.078 <0.01<0.020.03 4.4 131.0 17.0 138.0 0.135<0.0070.337 305 $32010.0<0.2 <0.004133.0<0.01<0.05 <0.01<0.020.04 4.7 124.0 5.0 190.0 <0.05<0.0070.334 444 $46711.0<0.2 <0.004128.0<0.01<0.05 <0.01<0.020.03 5.6 97.0 0.1 243.0 <0.05<0.0070.318 589 $616*The bolded row represents the point atwhich the water approximately meets Pennsylvania effluent standards.** Assumes the cost of a 20% NaOH solution is $0.50/gallonTable 10. Results of treatment acidity titration for Manor.Dissolved metal concentrations as a function of pH.TitrationpHendpointAl As Ca Cd Co CuFe3+Fe2+K Mg Mn Na Ni Se ZnTreatmentNetAciditymg/LCaCO3AnnualCost totreat aflow of 1gpm**2.98 (raw)33.0 <0.004 190.0<0.01 0.18 0.03 0.03326.0 2.8 57.0 4.7 3.8 0.294<0.0071.2855.60.25 <0.004--<0.010.29 <0.01<0.02326.0 6.6 -- -- -- 0.508<0.0071.824 369 $3887.0<0.2 <0.004172<0.010.124 <0.01<0.02116.0 3.8 49 3.9 269.0 0.125<0.0070.486 678 $7139.0<0.2 <0.004172<0.010.242 <0.01<0.022.5 3.1 49 1.2 374.0 <0.05<0.0070.397 864 $90910.00.32 <0.004203<0.010.177 <0.01<0.021.0 3.9 52 0.1 482.0 <0.05<0.0070.412 929 $97711.00.05 <0.004155<0.010.078 <0.01<0.020.1 3.4 19 0.03 438.0 <0.05<0.0070.352 1024 $1,077*The bolded row represents the point at which the water meets Pennsylvania effluent standards.** Assumes the cost of a 20% NaOH solution is $0.50/gallon1263Table 11. Results of treatment acidity titration for Millerstein. Dissolved metal concentrations as a function of pH.TitrationpHendpointAl As Ca Cd Co CuFe3+Fe2+K Mg Mn Na Ni Se ZnTreatmentNetAciditymg/LCaCO3AnnualCost totreat aflow of 1gpm**5.98 (raw)0.25 <0.004 101.0<0.01 <0.05<0.01<0.024.12 2.8 113.0 5.2 46.0 0.11<0.0070.1057.5<0.2 <0.004101.0<0.01<0.05<0.01<0.02<0.023.0 113.0 4.9 55.0 0.07<0.0070.332 19 $208.9<0.2 <0.004102.0<0.01<0.05<0.01<0.02<0.023.6 115.0 4.9 60.0 0.06<0.0070.296 26 $279.5<0.2 <0.00499.0<0.01<0.05<0.01<0.02<0.023.0 110.0 3.7 63.0 <0.05<0.0070.255 40 $4210.5<0.2 <0.004105.0<0.01<0.05<0.01<0.020.05 3.3 111.0 0.019 92.0 <0.05<0.0070.284 96 $10111.0<0.2 <0.004102.0<0.01<0.05<0.01<0.020.036 3.1 107.0 0.013 101.0 <0.05<0.0070.272 135 $142*The bolded row represents the point at whichthe water meets Pennsylvania effluent standards.** Assumes the cost of a 20% NaOH solution is $0.50/gallonTable 12. Results of treatment acidity titration for Morris. Dissolved metal concentrations as a function of pH.TitrationpHendpointAl As Ca Cd Co CuFe3+Fe2+K Mg Mn Na Ni Se ZnTreatmentNetAciditymg/LCaCO3AnnualCost totreat aflow of1 gpm**3.01 (raw)61.0 <0.004 186.0<0.011.3300.033 <0.020.75 3.8 460.0 57.0 18.0 2.1 0.0074 4.1224.727.0 <0.004189.0<0.011.321 0.098<0.020.07 4.1 460.0 56.0 113.0 2.06 0.0086 4.066 226.0 $2375.8<0.2 <0.004188.0<0.011.117 <0.01<0.02<0.024.0 460.0 52.0 185.0 1.76 0.0071 2.168 401.0 $4217.0<0.2 <0.004184.0<0.010.945 <0.01<0.02<0.023.8 452.0 50.0 0.2 1.57 0.0079 1.124 432.0 $4548.30.5 <0.004191.0<0.010.416 <0.01<0.02<0.023.9 457.0 41.0 229.0 0.78 0.0093 0.596 506.0 $5329.00.7 <0.004174.0<0.010.133 <0.01<0.02<0.023.8 442.0 30.0 263.0 0.21 0.0089 0.431 595.0 $6259.50.91 <0.004174.0<0.010.052 <0.010.034<0.024.2 430.0 19.0 280.0 <0.05 0.0089 0.407 619.0 $65010.00.97 <0.004173.0<0.01<0.05 <0.010.033<0.024.1 412.0 4.6 349.0 <0.05 0.0079 0.398 749.0 $78611.0<0.2 <0.004 173.0<0.01<0.05 <0.010.026 <0.026.0 194.0 0.1 726.0 <0.05 0.0112 0.373 1743.0 $1,827*The bolded row represents the point at which the water meets Pennsylvaniaeffluent standards.** Assumes the cost of a 20% NaOH solution is $0.50/gallon1264 Table 13. Acidity determined by each of the methods.

Calculated Net Acidity Hot Acidity pH 8.2 Hot Acidity pH
11.0 Treatment Net Acidity
Ike 120.4 123 720 444
Manor 845.0 880 1360 864
Millerstein 4.7 16 1340 19
Morris 549.4 510 2840 749 Acidity expressed as mg/L of CaCO 3 . Identifying the sources of acid ity measured by the hot acidity (pH 8.2) titrations is not as straight forward for Manor, Morr is Millerstein as it was for Ike since differences between
calculated net acidities and hot acidities exist. Therefore, the acidit y produced during the hot acidity titrations was modeled by using difference in dissolved metal concentrations between the
raw water and the end of the hot acidity titration at pH 8.2 and 11.0 as input into Equation (9). The amount of metal precipitation that occurred during the Hot Ac idity titration to pH 8.2 and pH 11.0 is presented in Tables 7 and 8 for each of the four discharges. Table 14 presents the results of using data from Tables 7 and 8 as input into Equation (9) to model the acidity produced during the hot acidity titrations, first considerin g only the pH and concentrations of Fe, Mn, and Al, and then also considering the concentration of Mg. The results of Table 13 and Table 14 should be identical if complete hydrolysis of Al, Fe, and Mn occurred during the hot acidity
titration. If the results in Table 14 are lower th an the calculated aciditi es in Table 13, then incomplete hydrolysis of Al, Fe, or Mn occu rred during the hot aci dity titration. Unlike Ike, hot acidity (pH 8.2) for the Morris water is 39.4 mg/L lower than calculated net acidity (Table 13). Likewise, the results of modeling the acidity produced from Al, Fe, and Mn during the hot acidity (pH 8.2) titration (Table 14 ) are lower than calculated net acidity in Table 13. Table 8 shows that 21.2 mg/L of Mn was s till dissolved in the Morris water at the end of the hot acidity titration. A concen tration of 21.2 mg/L of dissolved Mn equa tes to 38.5 mg/L of acidity. The residual Mn in solution at the end of the hot acidity titration (pH 8.2) explains why the hot acidity (pH 8.2) is lower than calculated acidity. The incomplete removal of Mn during the hot acidity titration is probably not a result of solubility constraints or insufficient addition of hydrogen peroxide. It is probably the result of the lab chemis t prematurely ending the titration. This result is important because it shows hot acidity titrations (pH 8.2) do not always measure the maximum contribution of acidity from Mn.1265 Table 14. Results of using Equations (9) and (12) and data in Tables 7 and 8 to model acidity produced during the hot acidity titrations [Acidity e xpressed as mg/L of CaCO 3 ]

Hot Acidity @ pH 8.2 Calculated Net Acidity @ pH 8.2 considering pH, Al, Fe, and

Mn acidity Calculated Net Acidity @ pH 8.2 considering pH, Al, Fe, Mn, and Mg acidity Hot Acidity @ pH 11.0 Calculated Net Acidity @ pH 11.0 considering pH, Al, Fe, and

Mn acidity Calculated Net Acidity @ pH 11.0 considering pH, Al, Fe, Mn, and Mg acidity

Ike 123 113 177 720 143 609
Manor 880 846 881 1360 849 1082
Millerstein 16 4.3 17.2 1340 4.3 338
Morris 510 502 722 2840 546 2496

Unlike Ike or Morris, Table 13 shows that hot acidity (pH 8.2) results for Manor and Millerstein are higher than the co rresponding calculated net acidity. Hot acidity (pH 8.2) is 35 mg/L higher than calculated net acidity for Manor and 11.3 mg/L higher for Millerstein. These results suggest that these hot acidity titrations are measuring another source of acidity not considered by Equation (9). At higher pH, ma gnesium and calcium can undergo hydrolysis and produce acidity, as explained below.

+ + + = + H OH Mg O H Mg 2 ) ( 2 2 2 (11)

+ + + = + H OH Ca O H Ca 2 ) ( 2 2 2 2 (12)

Empirical data from this study (Tables 7 & 8, Tables 9-12) a nd the authors’ experience show Mg hydrolysis reactions are signific ant at pH > 10.0 and Ca solubility is not likely to be affected by hydrolysis reactions until pH > 11.0. The effect of the Mg hydrolysis process on Mg solubility is illustrated in Fig. 1. Fig. 1 show s, in theory, the concentration of Mg can decline over five orders of magnitude between th e two hot acidity pH endpoints, from 10 0.44 mol/L (66,980 mg/L) at pH of 8.2 to 10 -5.0 mol/L (0.23 mg/L) at pH of 11. This solubility is consistent with that observed during some of the hot acid ity titrations. Small amounts of Mg precipitation occurred during hot acidity titrations to pH 8. 2 and large amounts of Mg precipitation occurred during hot acidity titration to pH 11.0 (Tables 7 & 8). Thus, E quation (13), which is consistent with Equation (8), can be used to model the acidity (mg/L of CaCO 3 ) resulting from Mg1266 hydrolysis at high pH. It should not be automa tically assumed that re ductions in Mg and Ca during acidity titrations resulted in acidity production. Reduction in Mg and Ca may be attributable to adsorption to precipitates or to the preci pitation of calcium carbonate.

              × × = + 24 2 50 2 Mg MgAcidity (13)

While the results show that hot acidities for Manor and Millerstein are greater than calculated net acidities, the results also show Mg precipitated during the hot acidity titrations. Tables 7 and 8 show that 8.5 mg/L of Mg prec ipitated during the Manor Hot Acid ity titration to pH 8.2 and 3.1 mg/L precipitated during the Millerstein titration. Equation (13) was used to evaluate whether the difference between the acidity measured by the hot acidity (pH 8.2) ti trations and calculated acidity for Manor and Millerstei n may be due to the hydrolysis of Mg. The hydrolysis of 8.5 mg/L of Mg results in 35 mg/L of acidity and h ydrolysis of 3.1 mg/L of Mg results in 12.9 mg/L of acidity. As Table 14 shows, the addition of the Mg acidity to the calculat ed net acidity that only considers Al, Fe, and Mn hydrolys is results in acidity values th at are almost identical to the hot acidities (pH 8.2) for Manor an d Millerstein. The hydrolysis of Mg is the source of acidity that resulted in a higher hot ac idity (pH 8.2) for Manor and Mill erstein. These results suggest that for Manor and Millerstein a hot acidity titra tion to pH 8.2 mainly measured acidity due to pH, Al, Fe, and Mn but also measured a small amount of acidity from the hydrolysis of Mg.
This is consistent with the work of Payne and Yeates (1970).12670123456789102 4 6 8 10 12 14pH

-log [M TOT (mol/L)]Al(OH)3Ca(OH)2Cd(OH)2Co(OH)2Cu(OH)2Fe(OH)2Fe(OH)3Mg(OH)2Mn(OH)2Ni(OH)2Zn(OH)2Figure 1. Comparison of metal hydroxide solubilities for constituentscommonly found in acidic minedrainage (C. A. Cravotta, III,2004, written commun.). Equilibrium constants from Ball and Nordstrom (1991) at 25oC.1268 A similar process to that used above was used to try to identify the sources of acidity measured during the treatment titra tions (Tables 9 – 12). It was assumed that a decline in metal concentration during the treatment titrations was a result of hydrol ysis reactions. The amount of metal precipitation (via hydrolysis ) that occurred between each tr eatment titration was calculated by subtracting the metal concentrations at the end of each treatment titration from the metal
concentrations in the ra w water (Tables 9 – 12). These results were then used as input into Equation (9) to calculate acidity production from pH, Al, Fe, and Mn. Tables 15 – 18 show the results of using Equation (9) to model the tr eatment titrations. The following discussion will
focus on the bolded rows in Tables 9-12 and 15- 18 since these rows represent the (approximate) point at which effluent standards are met. Tables 15-18 show that for all cases but Millerstein, the treatment titration acidity is considerably high er than the results using Equation (9). This is probably because only a small am ount of Mn had to be precipita ted for Millerstein to meet effluent standards, so the pH did not have to be raised as high as normally required for significant Mn removal. Tables 15-18 show that it is at high pH (e.g. pH > 9) that sources of acidity not considered by Equation (9) become si gnificant (i.e. treatment acidities > calculated net acidities). The other waters ha d either high leve ls of Mn or Fe 2+ , which require high pH for precipitation by hydrolysis. Ike, Manor, and Morris did not meet e ffluent standards until the pH was raised to approximately 10.0, 9.0, and 11.0, respectively. As Tables 15, 16, and 18 show, the large difference between the calculated net acidity and the treatment titration acidity indicates that another source(s) of acidity , not considered by E quation (9) was present which resulted in much higher treatment titration acidities. Tables 9 – 12 show large reduc tions in Mg occurred during the Treatment Titrations at high pH. While the results of the Manor and Millerstein hot acidity titrations showed that Mg produced acidity during the hot acidity titrations, the results of Fig. 2 and Table 11 provide evidence that Mg pr oduced acidity during the treatment titrations. Fig. 2 and Table 11 show that as the pH of the Morris water is raised from 10.0 to 11.0, treatment net acidity raises from 749 to 1743 mg/L (as CaCO 3 ), while the Mg concentration decreases by 218 mg/L. Therefore, Equations (9 ) and (13) were combined to evaluate whether Mg acidity1269Table 15. Results of using Equations (9) and(12) and data in Table 9to model acidity producedacidity produced during the treatment acidity titrations for Ike [Acidity expressed as mg/L of CaCO3]IkeTitration pHTreatment NetAcidityCalculated Net Acidity @ considering pH,Al, Fe, and Mn acidityCalculated Net Acidity considering pH, Al, Fe, Mn, andMg acidity4.18 (raw)Raw5.539 39 396.570 54 547.0207 76 858.0232 82 999.0305 90 10210.0444 111 15311.0589 121 275Table 16. Results of using Equations (9) and (12) and data in Table 10 to model acidity producedduring the treatment acidity titrations for Manor [Acidity expressed as mg/L of CaCO3]ManorTitration pHTreatment NetAcidity(mg/L CaCO3)Calculated Net Acidity @ considering pH,Al, Fe, and Mn acidityCalculated Net Acidity considering pH, Al, Fe, Mn, and Mg acidity2.98 (raw)Raw5.6369 182 1827.0678 611 6449.0864 818 85110.0929 822 84311.01024 821 9791270Table 17. Results of using Equations (9)and (12) and data in Table 11to model acidity producedduring the treatment acidity titrationsfor Millerstein [Acidity expressed as mg/L of CaCO3]MillersteinTitration pHTreatment NetAcidityCalculated Net Acidity @ considering pH,Al, Fe, and Mn acidityCalculated Net Acidity considering pH, Al, Fe, Mn, and Mg acidity5.98 (raw)Raw7.519 7 78.926 7 79.540 9 2210.596 16 2411.0135 16 41Table 18. Results of using Equations (9) and (12) and data inTable 12 to model acidity produced during the treatment aciditytitrations for Morris [Acidity expressed as mg/L of CaCO3]MorrisTitration pHTreatment NetAcidityCalculated Net Acidity @considering pH, Al, Fe, and MnacidityCalculated Net Acidity consideringpH, Al, Fe, Mn, and Mg acidityCalculated Net Acidity considering pH, Al, Fe,Mn, Mg, Mg(OH)+, Ca(OH)+, CaHSO4+, andNaOH (aq) acidity3.01 (raw)Raw4.7226.0 216 2165.8401.0 373 3737.0432.0 376 4108.3506.0 391 21339.0595.0 410 4859.5619.0 404 52910.0749.0 430 63011.01743.0 442 1550 15901271Figure 2. Results of treatment titration for the Morris water. Note the relationship between the sharp increase in acidity andthe sharpdecrease in Mg concentration at pH 11.0.Dissolved Metal concentration and Acidity vs. pH for Morris0102030405060702.50 4.50 6.50 8.50 10.50pH

Al, Fe and Mn (mg/L)03006009001200150018002100 Mg and Acidity (mg/L)AlFe2+MnMgTreatment Acidity1272 may explain the large difference between treatment acidities and calculated net acidities at high pH. The results in Tables 15 – 18 show the treatment titration acidities are still considerably higher than calculated acidities, even when considering Mg hydrolys is. This result suggests that additional sources of acidity, not considered by Equations (9) or (12), are being measured by the treatment titrations, especi ally at high pH. Other possible sour ces of acidity are the formation of cation-hydroxyl complexes at high pH . Waters like Morris that ha ve large amount of Mg and Ca left in solution at high pH can consume hydroxyl to form Mg(OH) + , Ca(OH) + , and CaHSO 4 + complexes. The base-consuming effect of th ese complexes may be the source of the unknown acidity. Other sources of acidity not accounted for by Equation (9) are bisulfate acidity and carbonate acidity. While identifying the exact s ources of acidity produ ced during all of the treatment titrations will be the focus of future work, the geoche mical modeling program PHREEQC (Parkhurst and Appelo, 1999) was used to dete rmine whether the formation of cation-hydroxyl complexes in the Morris water at pH 11.0 could account for the additional acidity measured by the treatment
titrations. Table 18 shows the di fference in acidity measured by the treatment titration at pH 11.0 and calculated acidity is 193 mg/L (as CaCO 3 ). Results of PHREEQC showed the baseconsuming complexes at pH 11.0 were equi valent to 40 mg/L of acidity (as CaCO 3 ), which does
not account for all of the unknown acidity. Further speciation calculations performed by PHREEQC show the raw water may have contained approximately 40 mg/L of bisulfate acidity. The authors hypothesize the remaining 113 mg/L of acidity may be from carbonate acidity (H 2 CO 3 ). Lime was purposely added to the mine fl oor during the reclamati on of the Morris site to aid in the prevention of acid mine drainage. The dissolution of the lime by acidic water would result in an increase of the part ial pressure of carbon dioxide (pCO 2 ). While the amount of CO 2 in the Morris water was not determined by laboratory analysis, the pCO 2 would have to be approximately 10 -1.42 to contribute approximately 113 mg/L of acidity (as CaCO 3 ).
The authors also believe that carbonate ac idity is the reason why acidity measured by treatment titrations is often higher than calculate d acidities at low to mid pH endpoints (Tables 15-18). Jageman et al . (1988) has shown that carbonic acid can be a considerable source of acidity in mine drainage. Si nce neither calculated acidity no r hot acidity c onsider carbonate acidity, treatment titrations should be used to estimate the total acid ity that will need neutralized when chemically treating acid mine dr ainage to moderate or high pH.1273 Practical Considerations This study shows that the ac idity of mine drainage depends on the method used for determination and the titration endpoint. This is an important point because it proves that the commonly used Standard Methods 2 310 hot acidity should not be us ed to universally define the acidity of mine drainage, as is commonly the case. The results of this stu dy show that hot acidity (pH 8.2) titrations mostly measure acidity fr om pH, Al, Fe, and Mn, which are the regulated parameters in mine drainage. While mine drai nage contains elevated levels of numerous different metals, Al, Fe, and Mn are the metals that are capable of generating considerable acidity within both the pH range of natural waters and the pH ra nge of effluent water emanating from limestone-based passive treatment systems. Therefore, hot acidity (pH 8.2) measurements are suitable for assessing the impact to stream s receiving untreated mine drainage and for predicting the non-carbonate acidity that will need neutralized when passively treating water
with limestone. This study also showed that hot acidity (p H 8.2) titrations and calculated net acidities primarily measure identical sources of acidity. Nevertheless, the authors caution against using Equation (9) as a unanimous substitution for a hot acidity titration. Low pH mine water containing high levels of sulfate is likely to contain considerable bi sulfate. This source of acidity would be measured by a hot acidity titration but not considered by Equation (9). Misapplication of Equation (9) to waters containi ng considerable bisulf ate would result in a mischaracterization of the acidity encountered in the pH range of natural waters. Standard Methods 2310 hot acidity titrati on and calculated acid ity (Equation (9)) are designed to measure specific sources of acidity . In fact, the Standard Methods procedure requires the addition of oxidants and boiling to en sure only the desired sources of acidity are measured. Treatment titrations indiscriminately measure acidity, includ ing carbonate acidity, as pH is raised. As the results show, a treatment titration is the most appropriate method for measuring the acidity neutralized by chemical tr eatment systems treating to high pH. The results of a treatment titration are useful to operators of active treatment systems. For example, Table
11 shows the Morris water needs to be treated to (approximately) pH 10.0 to meet effluent standards. The consequence of over treating this water to pH 11.0 is severe. The acidity increases from 749 mg/L to 1743 mg/L as CaCO 3 and 219 mg/L of Mg precipitates. The effect of increased acidity and Mg precipitation is twofol d. First, the annual cost to treat each gallon1274 per minute of flow with a 20% solution of caustic soda increa ses from $786 to $1,827 (assuming $0.50/gal) (Table 12). Secondly, as Fig. 3 shows, the precipitati on of an additional 219 mg/L of Mg results in a doubling in sludge volume. On th e other hand, Table 10 shows that over treating the Manor water by two pH units w ill not double chemical costs, as is the case with Morris, since little Mg exists in the Manor water. Treatment ti trations provide treatment plant operators with
the information needed to increase treatment efficiency and reduce treatment costs.

Fi gure 3. Sludge volumes for Morri s discharge titration. Left to right, titrated to pH 9.0, titrated to pH 10.0, titrated to pH 11.0

An additional point worth men tioning is that every 1 mg/L of Mg is capable of producing 4.1 mg/L of acidity (as CaCO 3 ). This should be of great importa nce to operators tr eating mine water to high pH, since Mg is generally present in large amounts in mine dr ainage. The authors’ experience is that mine drainage emanating from surface mines in Pennsylvania and West Virginia generally has greater than 80 mg/L of Mg wh ereas mine drainage emanating from underground mines generally has less than 80 mg/L of Mg. Over treating mine drainage may release large amounts of Mg acidity.

Conclusion

This study showed that hot acidity measurements should not be used to universally define the acidity of mine drainage. The acidity of mi ne drainage varies with the method used for 

### pH 9.0 pH 11.0 pH 10.01275 determination and with the titr ation endpoint. For the most pa rt, this study showed Standard Methods 2310 hot acidity procedure and calcula ted acidity (Equation (9)) measure the same sources of acidity. These met hods are suitable for estimating the noncarbonate (i.e. pH and metal) acidity encountered while using a limestone -based passive treatment system to treat mine drainage, but are not suitable for estimating the acidity encountered while chemically treating water at high pH. The results of the treatment ti trations showed that la rge amounts of acidity are released at high pH. The data shows that the hydr olysis of Mg contribute s to acidity at high pH; however, additional unknown sources of acidity exist and their identification will be the focus of future work. The authors hypothesize that the formation of cation-hydroxyl complexes may be responsible for the additional acid ity. Data produced by the treatm ent titrations show that the behavior of Mg in mine water treated to hi gh pH increases both chem ical costs and sludge
volumes.

Acknowledgments

The authors are grateful to Dr. Art Rose of Penn State Univers ity, Dr. Chuck Cravotta of the U.S. Geological Survey, and Dr. Carl Kirby of Bucknell University for their valuable insight and
countless discussions on th e subject of acidity.

Literature Cited

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