1249

COMPARISON OF THREE METHODS TO MEASURE ACIDITY OF 

COAL-MINE DRAINAGE1 

Brent Means2 and Tiff Hilton

Abstract. Although the Standard Methods 2310 hot peroxide acidity procedure is widely 

used for measuring the acidity of mine drainage, 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 methods for estimating the acidity of net-acidic 

waters emanating from the Manor, Millerstein, Ike, and Morris coal mines in 

Pennsylvania: Standard Methods 2310 hot acidity 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 acidity 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 resulted from incomplete hydrolysis of Mn 

during the hot acidity titration. The agreement between the acidity measured by the 

treatment acidity titration and the other two acidity methods was within 16 mg/L (as 

CaCO3) for the Manor and Millerstein waters, but greater than 200 mg/L (as CaCO3) for 

the Ike and Morris waters. The fair agreement 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 formation 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 base-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 drainage. 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 25th 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 Hilton, Mining Engineer, WOPEC 

Inc., RR2 Box 294B, Lewisburg, WV 24901 (wopec@charter.net, Tel. 304-645-7633). 
1250

Introduction

It has long been recognized that acidity is an important measure for characterizing mine 

drainage (Rose and Cravotta, 1998; ADTI, 2000; Kirby, 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 acidity. Both cold and hot acidity (A.P.H.A, 1998) are laboratory 

procedures that measure acidity by titrating water 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-mining professionals in Appalachia have long 

considered a Standard Methods 2310 hot acidity titration 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 result of regulatory agencies, such as Pennsylvania Department of 

Environmental Protection (PADEP), requiring that acidity measurements (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 (Rose and Cravotta, 1998) and to calculate 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. However, little work has been 

done to describe the acidity of mine drainage at titration endpoints 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 three acidity methods used in this study. 

Second, a site selection section discusses the reasoning behind choosing the four mine waters. 

Third, a discussion section presents the results of the applying the acidity methods to the four 

mine waters. Finally, practical implications of the results are discussed. 
1251

Background

Overview of Acidity in Mine Drainage

Mine drainage commonly contains elevated amounts 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 Appalachian coal fields have aluminum, iron, or 

manganese concentration of greater than 50 mg/L (unpublished data). At favorable pH 

conditions, Al3+, Fe2+, Fe3+, and Mn2+ metals participate in hydrolysis reactions and produce 

acidity: 

 + + Al + 3H2 0 = Al(OH)3 + 3H 3 (1) 

 + + Fe + 0.25O2 + 2.5H2O = Fe(OH)3(s) + 2H 2 (2) 

 + + Fe + 3H2 0 = Fe(OH)3 + 3H 3 (3) 

 + + Mn + 0.5O2 + H2O = MnO2 (s) + 2H 2 (4) 

Fe2+and Mn2+ are redox dependent and are shown in equations (2) and (4) as precipitating as 

chemically oxidized solids, which are typically formed 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, Fe2+ and Mn2+ can precipitate as 

chemically reduced solids (Equations (5) and (6)); however, the acidity generated is identical to 

that expressed in Equations (3) and (4). 

 + + Fe + 2H2 0 = Fe(OH) 2 + 2H 2 (5) 

 + + Mn + 2H2 0 = Mn(OH) 2 + 2H 2 (6) 

In addition to elevated metals, many Appalachian discharges are characterized by low pH. 

OSM’s mine drainage inventory shows that 430 of 1470 post-primacy discharges in the 
1252

Appalachian coal fields have pH < 3.5 (unpublished data). Acidity from low pH (e.g. pH<3.0) is 

substantial and can be calculated by Equation (7): 

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

In addition to metal acidity and pH acidity, bisulfate (HSO4

-

), carbonic acid (H2CO3), and 

hydrogen sulfide (H2S) are known to contribute to the acidity of fresh mine drainage (Jageman et 

al., 1988; Rose and Cravotta, 1998). Payne and 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 expressed as mole equivalents. For 

instance, Equation (1) shows that 3 moles of acidity are produced for every 1 mole of Al3+

participating in a hydrolysis reaction. This stoichiometry permits the development of an 

equation that uses dissolved aluminum concentration to estimate acidity generation from 

hydrolysis and precipitation. 

 













 × = ×

+

+

27

3 50

3

3 Al Al acidity (8) 

where Al3+ Acidity is the amount of acidity in mg/L of CaCO3, Al3+ is dissolved Aluminum 

concentration in mg/L, and the values of 27 and 50 are the molecular weight of aluminum and 

the conversion factor for expressing acidity as CaCO3, 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. 

 



 





+ × 













 ×

+













 ×

+













 ×

+













 × = × −

+ + + + Fe Fe Al Mn pH Acidity 1000 10

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 CaCO3 and Fe2+, Fe 3+, Al3+, 

and Mn2+ are the analytical concentrations of dissolved metals in mg/L. 

 Equation (9) uses results 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 sensitive 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 CaCO3) and 

100 mg/L (as CaCO3), respectively. This example illustrates the importance of using a calibrated 

pH meter to measure field pH when using Equation (9). A laboratory-measured pH should not 

be substituted for a field-measured pH since a lab pH could be unrepresentative if pH-altering 

chemical reactions, such as hydrolysis and exsolution of CO2, 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 chemical species contained in a field-filtered (0.45 

μm) sample as dissolved, while an unfiltered sample 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 preservation 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 Al3+ (mg/L) 

Calculated Acidity from Al hydrolysis 

(mg/L CaCO3) 

Sample Filtered 5 28 

Sample NOT Filtered 10 56 

 For the purposes of comparing the results of the three acidity methods, net acidity was 

calculated by subtracting alkalinity from calculated acidity (Equation 9): 

 Net Calculated Acidity (mg/L CaCO3) = 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 specified pH endpoint. Several published hot acidity 

procedures exist (APHA, 1998; ASTM, 1984; U.S.EPA, 1979), however, most laboratories in 

Pennsylvania use either Standard Methods 2310 (APHA, 1998) or EPA 305.1 (U.S. EPA, 1979) 

procedures (Means, unpublished data). Although 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 three methods employ consistent steps. Kirby and Cravotta (2004) 

offer a thorough discussion of the hot acidity titration procedure, resulting in an abbreviated 

discussion for this paper. The general hot acidity steps are to (1) titrate the water using a sulfuric 

acid to pH less than 4.5 to neutralize alkalinity; (2) add five drops of a 30% solution of hydrogen 

peroxide (H2O2) for metal oxidation; (3) boil the sample for two to five minutes to remove all 

carbonate species, accelerate oxidation and precipitation of metals, and then allow the sample to 

cool to room temperature; (4) titrate the water using a strong base to pH 8.3 to neutralize acidity; 

and (5) calculate hot acidity by subtracting the equivalent of acid added from the equivalent of 

base added and express hot acidity as mg/L of CaCO3. For the purpose of this paper, there are 

two important points in this procedure. First, the procedure intentionally does not measure 

carbonate acidity (H2CO3). Secondly, step (5) yields 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 water to effluent limits. A treatment acidity 

titration is a variant of a “conventional” cold acidity titration. There are two general differences 

between a treatment titration and a cold acidity titration. First, unlike a conventional cold acidity 

titration, a treatment acidity titration is a group of individual titrations with each titration using a 

different pH endpoint. The person performing the titration determines 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 increased 
1255

and provides the pH where effluent standards are met. The second difference is the point at 

which the titration ends. In most cases, a “conventional” cold acidity titration 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 reached and pH is free to drift downwards as 

hydrolysis reactions continue to occur. This distinction stems from the difference in goals for 

each procedure. The goal of a cold acidity titration is to calculate 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 chemical 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 measured from both methodologies will differ. 

 Table 2. These pH ranges do not necessarily correspond to the minimum solubility for 

each species, but rather correspond to the approximate ranges in pH needed to meet 

effluent standards for a specific metal (based on authors’ experience). 

Metal pH range needed to meet metal effluent 

limits 

Fe3+ 3.5 – 4.5 

Al3+ 4.5-6.0 

Fe2+ 7.0 – 9.0 

Mn2+ 9.0 – 10.5 

 The basic procedure for a treatment titration is as follows: (1) Pour 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 titrate the water to the target pH endpoint. 

Record the amount of NaOH consumed. (3) Filter approximately 125 ml of the titrated water 

using a hand pump and a 0.45-μm filter. Add HNO3 until the pH of the sample is less than 2.0 to 

preserve the sample for laboratory metal analysis. (4) Repeat Steps 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 laboratory analysis of metals concentration to 

determine the point at which the water meets effluent limits and calculate the treatment acidity. 
1256

Water Quality Collection and Analysis

Calculated Acidity

 The following procedure was used to determine calculated acidities. In the field, mine 

drainage was passed through a 0.45-μm 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 HNO3 to pH < 2.0. Cardan 

Laboratories of Northern Cambria, Pennsylvania determined metal 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 and 7.0 buffer solutions. A sample was collected, placed on ice, 

and analyzed for alkalinity within 24 hours using a Hach alkalinity titration kit. 

Hot Acidity

 Unfiltered water samples were collected and placed on ice until hot acidity titrations were 

performed. During the manual hot acidity titrations, Cardan Laboratories 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 the 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 refrigerated and treatment acidity titrations were 

performed within 24 hours. Pennsylvania Bureau of Laboratories determined metal 

concentrations from the filtered water obtained at each of the different titration endpoints using 

inductively coupled plasma emission 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 desirable to study different sources of acidity (Tables 3 and 4). The 

discharges Ike, Millerstein, and Morris are from reclaimed surface mines and Manor emanates 

from an underground mine. Ike was selected because of its moderate metal 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 concentration 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 Fe3+ Fe2+ K Mg Mn Na SO4

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 data 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 4.122 
1258

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 various 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 column presents the dissolved concentrations after the hot peroxide 

boil. The third and fourth columns show the dissolved 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 numerous pH endpoints, filtered, and analyzed for 

dissolved metals. These tables also show the corresponding acidity for each incremental 

titration. The bolded rows in these tables represent the titration that resulted in (approximately) 

meeting effluent standards defined by Pennsylvania regulations. Table 13 provides a summary 

of the results for each acidity method. 

Discussion of Results

 Table 13 shows that the hot acidity titrations performed to the endpoint pH of 11.0 produced 

the highest acidity values, followed by the treatment 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 methods that produced different results are measuring different 

sources of acidity. In order to correctly apply each acidity method in the future, 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 hydrolysis of Al, Fe, and Mn. Table 13 shows that 

calculated net acidities and hot acidities (pH 8.2) for Ike are virtually identical, which suggests 

both methods measure the same sources of acidity. 
1259

Table 5. Raw water analyses performed 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 -- -- 

SO4

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 CaCO3) 

Field Alkalinity 

(mg/L CaCO3) 

Net Acidity 

(mg/L CaCO3) 

Ike 122.4 2.0 120.4 

Manor 845.0 0.0 845.0 

Millerstein 20.7 16.0 4.7 

Morris 549.4 0.0 549.4 
1260

Table 7. Hot Acidity titration for the Ike and Manor 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 

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 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 

SO4

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 CaCO3 
1261

Table 8. Hot Acidity titration for the Millerstein 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 CaCO3
1262

Table 9. Results of treatment acidity titration for Ike. Dissolved metal concentrations as a function 

Titration pH endpoint 

Al As Ca Cd Co Cu Fe3+ Fe2+ K Mg Mn Na Ni Se Zn 

Tr

AC

4.18 (raw) 5.95 <0.004 150.0 <0.01 0.34 0.01 <0.02 24.0 4.5 134.00 24.0 7.5 0.48 <0.007 1.032 5.5 1.0 <0.004 150.0 <0.01 0.314 <0.01 <0.02 19.0 4.5 138.0 24.0 25.0 0.485 <0.007 1.068 6.5 0.2 <0.004 150.0 <0.01 0.289 <0.01 <0.02 13.0 5.3 138.0 24.0 40.0 0.454 <0.007 0.936 7.0 <0.2 <0.004 144.0 <0.01 0.242 <0.01 <0.02 1.6 4.9 132.0 23.0 95.0 0.343 <0.007 0.486 8.0 <0.2 <0.004 141.0 <0.01 0.177 <0.01 <0.02 0.42 4.4 130.0 21.0 113.0 0.281 <0.007 0.435 9.0 <0.2 <0.004 141.0 <0.01 0.078 <0.01 <0.02 0.03 4.4 131.0 17.0 138.0 0.135 <0.007 0.337 10.0 <0.2 <0.004 133.0 <0.01 <0.05 <0.01 <0.02 0.04 4.7 124.0 5.0 190.0 <0.05 <0.007 0.334 11.0 <0.2 <0.004 128.0 <0.01 <0.05 <0.01 <0.02 0.03 5.6 97.0 0.1 243.0 <0.05 <0.007 0.318 

 *The bolded row represents the point at which the water approximately meets Pennsylvania effluent standards. 

 ** Assumes the cost of a 20% NaOH solution is $0.50/gallon 

Table 10. Results of treatment acidity titration for Manor. Dissolved metal concentrations as a funct

Titration 

pH 

endpoint 

Al As Ca Cd Co Cu Fe3+ Fe2+ K Mg Mn Na Ni Se Zn 

2.98 (raw) 33.0 <0.004 190.0 <0.01 0.18 0.03 0.03 326.0 2.8 57.0 4.7 3.8 0.294 <0.007 1.285 

5.6 0.25 <0.004 -- <0.01 0.29 <0.01 <0.02 326.0 6.6 -- -- -- 0.508 <0.007 1.824 7.0 <0.2 <0.004 172 <0.01 0.124 <0.01 <0.02 116.0 3.8 49 3.9 269.0 0.125 <0.007 0.486 9.0 <0.2 <0.004 172 <0.01 0.242 <0.01 <0.02 2.5 3.1 49 1.2 374.0 <0.05 <0.007 0.397 

10.0 0.32 <0.004 203 <0.01 0.177 <0.01 <0.02 1.0 3.9 52 0.1 482.0 <0.05 <0.007 0.412 11.0 0.05 <0.004 155 <0.01 0.078 <0.01 <0.02 0.1 3.4 19 0.03 438.0 <0.05 <0.007 0.352 

 *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/gallon 
1263

Table 11. Results of treatment acidity titration for Millerstein. Dissolved metal concentrations as a 

Titration 

pH 

endpoint 

Al As Ca Cd Co Cu Fe3+ Fe2+ K Mg Mn Na Ni Se Zn 

5.98 (raw) 0.25 <0.004 101.0 <0.01 <0.05 <0.01 <0.02 4.12 2.8 113.0 5.2 46.0 0.11 <0.007 0.105 

7.5 <0.2 <0.004 101.0 <0.01 <0.05 <0.01 <0.02 <0.02 3.0 113.0 4.9 55.0 0.07 <0.007 0.332 8.9 <0.2 <0.004 102.0 <0.01 <0.05 <0.01 <0.02 <0.02 3.6 115.0 4.9 60.0 0.06 <0.007 0.296 9.5 <0.2 <0.004 99.0 <0.01 <0.05 <0.01 <0.02 <0.02 3.0 110.0 3.7 63.0 <0.05 <0.007 0.255 

10.5 <0.2 <0.004 105.0 <0.01 <0.05 <0.01 <0.02 0.05 3.3 111.0 0.019 92.0 <0.05 <0.007 0.284 11.0 <0.2 <0.004 102.0 <0.01 <0.05 <0.01 <0.02 0.036 3.1 107.0 0.013 101.0 <0.05 <0.007 0.272 

 *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/gallon 

Table 12. Results of treatment acidity titration for Morris. Dissolved metal concentrations as a funct

Titration 

pH 

endpoint 

Al As Ca Cd Co Cu Fe3+ Fe2+ K Mg Mn Na Ni Se Zn 

3.01 (raw) 61.0 <0.004 186.0 <0.01 1.330 0.033 <0.02 0.75 3.8 460.0 57.0 18.0 2.1 0.0074 4.122 

4.7 27.0 <0.004 189.0 <0.01 1.321 0.098 <0.02 0.07 4.1 460.0 56.0 113.0 2.06 0.0086 4.066 5.8 <0.2 <0.004 188.0 <0.01 1.117 <0.01 <0.02 <0.02 4.0 460.0 52.0 185.0 1.76 0.0071 2.168 7.0 <0.2 <0.004 184.0 <0.01 0.945 <0.01 <0.02 <0.02 3.8 452.0 50.0 0.2 1.57 0.0079 1.124 8.3 0.5 <0.004 191.0 <0.01 0.416 <0.01 <0.02 <0.02 3.9 457.0 41.0 229.0 0.78 0.0093 0.596 9.0 0.7 <0.004 174.0 <0.01 0.133 <0.01 <0.02 <0.02 3.8 442.0 30.0 263.0 0.21 0.0089 0.431 9.5 0.91 <0.004 174.0 <0.01 0.052 <0.01 0.034 <0.02 4.2 430.0 19.0 280.0 <0.05 0.0089 0.407 

10.0 0.97 <0.004 173.0 <0.01 <0.05 <0.01 0.033 <0.02 4.1 412.0 4.6 349.0 <0.05 0.0079 0.398 11.0 <0.2 <0.004 173.0 <0.01 <0.05 <0.01 0.026 <0.02 6.0 194.0 0.1 726.0 <0.05 0.0112 0.373 

*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/gallon
1264

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 CaCO3. 

 Identifying the sources of acidity measured by the hot acidity (pH 8.2) titrations is not as 

straight forward for Manor, Morris Millerstein as it was for Ike since differences between 

calculated net acidities and hot acidities exist. Therefore, the acidity 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 Acidity 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 considering 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 than the calculated acidities in Table 13, then 

incomplete hydrolysis of Al, Fe, or Mn occurred during the hot acidity 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 still dissolved in the Morris water at the end of the 

hot acidity titration. A concentration of 21.2 mg/L of dissolved Mn equates 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 chemist 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 expressed as mg/L of CaCO3] 

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 corresponding 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, magnesium and calcium can undergo hydrolysis and 

produce acidity, as explained below. 

 + + Mg + H2O = Mg(OH) 2 + 2H 2 (11) 

 + + Ca + 2H2O = Ca(OH) 2 + 2H 2 (12) 

Empirical data from this study (Tables 7 & 8, Tables 9-12) and the authors’ experience show 

Mg hydrolysis reactions are significant 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 shows, in theory, the concentration of Mg can decline 

over five orders of magnitude between the two hot acidity pH endpoints, from 100.44 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 acidity 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, Equation (13), which is consistent 

with Equation (8), can be used to model the acidity (mg/L of CaCO3) resulting from Mg 
1266

hydrolysis at high pH. It should not be automatically assumed that reductions 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 precipitation 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 precipitated during the Manor Hot Acidity 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) titrations and calculated 

acidity for Manor and Millerstein 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 hydrolysis 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 calculated net acidity that 

only considers Al, Fe, and Mn hydrolysis results in acidity values that are almost identical to the 

hot acidities (pH 8.2) for Manor and Millerstein. The hydrolysis of Mg is the source of acidity 

that resulted in a higher hot acidity (pH 8.2) for Manor and Millerstein. These results suggest 

that for Manor and Millerstein a hot acidity titration 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). 
1267

0123456789

10

2 4 6 8 10 12 14

pH

-log [MTOT (mol/L)]

Figure 1. Comparison of metal hydroxide solubilities for constituents commonly found in acidic mine drainag

2004, written commun.). Equilibrium constants from Ball and Nordstrom (1991) at 25 oC. 
1268

 A similar process to that used above was used to try to identify the sources of acidity 

measured during the treatment titrations (Tables 9 – 12). It was assumed that a decline in metal 

concentration during the treatment titrations was a result of hydrolysis reactions. The amount of 

metal precipitation (via hydrolysis) that occurred between each treatment titration was calculated 

by subtracting the metal concentrations at the end of each treatment titration from the metal 

concentrations in the raw 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 treatment 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 higher than the results using Equation (9). This is 

probably because only a small amount of Mn had to be precipitated 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 significant (i.e. treatment acidities > calculated 

net acidities). The other waters had either high levels of Mn or Fe2+, which require high pH for 

precipitation by hydrolysis. Ike, Manor, and Morris did not meet effluent 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 Equation (9) was present which resulted in 

much higher treatment titration acidities. Tables 9 – 12 show large reductions 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 produced 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 CaCO3), while the Mg concentration decreases by 

218 mg/L. Therefore, Equations (9) and (13) were combined to evaluate whether Mg acidity 
1269

 Table 15. Results of using Equations (9) and (12) and data in Table 9 to model acidity produce acidity produced during the treatment acidity titrations for Ike [Acidity expressed as mg/L of C

 Ike

Titration pH 

Treatment Net 

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

Mg acidity 

4.18 (raw) Raw 

5.5 39 39 39 6.5 70 54 54 7.0 207 76 85 8.0 232 82 99 9.0 305 90 102 

10.0 444 111 153 11.0 589 121 275 

 Table 16. Results of using Equations (9) and (12) and data in Table 10 to model acidity produced  during the treatment acidity titrations for Manor [Acidity expressed as mg/L of CaCO3] 

Manor

Titration pH 

Treatment Net 

Acidity 

(mg/L CaCO3) 

Calculated Net Acidity @ considering pH, 

Al, Fe, and Mn acidity Calculated Net Acidity considering pH, Al, Fe, Mn, an

2.98 (raw) Raw 

5.6 369 182 182 7.0 678 611 644 9.0 864 818 851 

10.0 929 822 843 11.0 1024 821 979 
1270

 Table 17. Results of using Equations (9) and (12) and data in Table 11 to model acidity produced  during the treatment acidity titrations for Millerstein [Acidity expressed as mg/L of CaCO3] 

Millerstein

Titration pH 

Treatment Net 

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

5.98 (raw) Raw 

7.5 19 7 7 8.9 26 7 7 9.5 40 9 22 

10.5 96 16 24 11.0 135 16 41 

Table 18. Results of using Equations (9) and (12) and data in Table 12 to model acidity produced during 

 titrations for Morris [Acidity expressed as mg/L of CaCO3] 

Morris

Titration pH 

Treatment Net 

Acidity 

Calculated Net Acidity @ 

considering pH, Al, Fe, and Mn 

acidity 

Calculated Net Acidity considering 

pH, Al, Fe, Mn, and Mg acidity 

Calculated Net Acidity c

Mn, Mg, Mg(OH)+, Ca

NaOH (aq

3.01 (raw) Raw 

4.7 226.0 216 216 5.8 401.0 373 373 7.0 432.0 376 410 8.3 506.0 391 2133 9.0 595.0 410 485 9.5 619.0 404 529 

10.0 749.0 430 630 11.0 1743.0 442 1550 1
1271

Figure 2. Results of treatment titration for the Morris water. Note the relationship between the sharp increase in 

decrease in Mg concentration at pH 11.0. 

Dissolved Metal concentration and Acidity vs. pH for Morris0

10203040506070

2.50 4.50 6.50 8.50 10.50

pH

Al, Fe and Mn (mg/L)

03006009001200150018002100

Mg and Acidity (mg/L) 
1272

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 hydrolysis. This result suggests that 

additional sources of acidity, not considered by Equations (9) or (12), are being measured by the 

treatment titrations, especially at high pH. Other possible sources of acidity are the formation of 

cation-hydroxyl complexes at high pH. Waters like Morris that have large amount of Mg and Ca 

left in solution at high pH can consume hydroxyl to form Mg(OH)+

, Ca(OH)+, and CaHSO4

+

complexes. The base-consuming effect of these 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 sources of acidity produced during all of the treatment titrations 

will be the focus of future work, the geochemical modeling program PHREEQC (Parkhurst and 

Appelo, 1999) was used to determine 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 difference in acidity measured by the treatment titration at pH 

11.0 and calculated acidity is 193 mg/L (as CaCO3). Results of PHREEQC showed the baseconsuming complexes at pH 11.0 were equivalent to 40 mg/L of acidity (as CaCO3), 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 

(H2CO3). Lime was purposely added to the mine floor during the reclamation 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 partial pressure of carbon dioxide (pCO2). While the amount of CO2

in the Morris water was not determined by laboratory analysis, the pCO2 would have to be 

approximately 10-1.42 to contribute approximately 113 mg/L of acidity (as CaCO3). 

 The authors also believe that carbonate acidity is the reason why acidity measured by 

treatment titrations is often higher than calculated 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. Since neither calculated acidity nor hot acidity consider carbonate 

acidity, treatment titrations should be used to estimate the total acidity that will need neutralized 

when chemically treating acid mine drainage to moderate or high pH. 
1273

Practical Considerations

This study shows that the acidity 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 2310 hot acidity should not be used to universally define the 

acidity of mine drainage, as is commonly the case. The results of this study show that hot acidity 

(pH 8.2) titrations mostly measure acidity from pH, Al, Fe, and Mn, which are the regulated 

parameters in mine drainage. While mine drainage 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 range of effluent water emanating 

from limestone-based passive treatment systems. Therefore, hot acidity (pH 8.2) measurements 

are suitable for assessing the impact to streams 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 (pH 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 bisulfate. This source of acidity 

would be measured by a hot acidity titration but not considered by Equation (9). Misapplication 

of Equation (9) to waters containing considerable bisulfate would result in a mischaracterization 

of the acidity encountered in the pH range of natural waters. 

 Standard Methods 2310 hot acidity titration and calculated acidity (Equation (9)) are 

designed to measure specific sources of acidity. In fact, the Standard Methods procedure 

requires the addition of oxidants and boiling to ensure only the desired sources of acidity are 

measured. Treatment titrations indiscriminately measure acidity, including 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 treatment 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 CaCO3 and 219 mg/L of Mg precipitates. The effect 

of increased acidity and Mg precipitation is twofold. First, the annual cost to treat each gallon 
1274

per minute of flow with a 20% solution of caustic soda increases from $786 to $1,827 (assuming 

$0.50/gal) (Table 12). Secondly, as Fig. 3 shows, the precipitation of an additional 219 mg/L of 

Mg results in a doubling in sludge volume. On the other hand, Table 10 shows that over treating 

the Manor water by two pH units will not double chemical costs, as is the case with Morris, since 

little Mg exists in the Manor water. Treatment titrations provide treatment plant operators with 

the information needed to increase treatment efficiency and reduce treatment costs. 

 Figure 3. Sludge volumes for Morris discharge titration. Left 

 to right, titrated to pH 9.0, titrated to pH 10.0, titrated to pH 11.0 

 

 An additional point worth mentioning is that every 1 mg/L of Mg is capable of producing 4.1 

mg/L of acidity (as CaCO3). This should be of great importance to operators treating mine water 

to high pH, since Mg is generally present in large amounts in mine drainage. 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 whereas 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 mine drainage varies with the method used for 

pH 9.0 pH 11.0 pH 10.0 
1275

determination and with the titration endpoint. For the most part, this study showed Standard 

Methods 2310 hot acidity procedure and calculated acidity (Equation (9)) measure the same 

sources of acidity. These methods 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 titrations showed that large amounts of acidity are 

released at high pH. The data shows that the hydrolysis of Mg contributes 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 acidity. Data produced by the treatment titrations show that the 

behavior of Mg in mine water treated to high pH increases both chemical costs and sludge 

volumes. 

Acknowledgments

The authors are grateful to Dr. Art Rose of Penn State University, Dr. Chuck Cravotta of the 

U.S. Geological Survey, and Dr. Carl Kirby of Bucknell University for their valuable insight and 

countless discussions on the subject of acidity. 

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The National Mine Land Reclamation Center, West Virginia University, Morgantown, WV, 

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American Public Health Association, 1998, Standard methods for the examination of water and 

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water: Philadelphia, v. 11.01, p. 125-133. 

Ball, J. W., and Nordstrom, D. K., 1991, User's manual for WATEQ4F with revised data base: 

U.S. Geological Survey Open-File Report 91-183, 189 p. 
1276

Cravotta, C. A., III, and Kirby, C. S., 2004, Acidity and Alkalinity in Mine Drainage: Practical 

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1277

Pennsylvania Department of Environmental Protection, 2004, Bituminous Surface Mine Permit 

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Prevention in Pennsylvania: Harrisburg, Pa., Pennsylvania Department of Environmental 

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