association quotients of aluminum sulphate complexes in nacl media from 50 to 125°c: results of a...

PII S0016-7037(99)00016-2

Association quotients of aluminum sulphate complexes in NaCl media from 50 to 125°C:Results of a potentiometric and solubility study

MOIRA K. RIDLEY,1,* DAVID J. WESOLOWSKI,2 DONALD A. PALMER,2 and RICHARD M. KETTLER1

1Department of Geology, University of Nebraska, Lincoln, NE 68588-0340, USA2Chemical and Analytical Sciences Division, Oak Ridge National Laboratory, P.O. Box 2008, Oak Ridge, TN 37831-6110, USA

(Received November5, 1997;accepted in revised form January8, 1998)

Abstract—The speciation and molal formation quotients for the complexation of aluminum with sulphatewere measured based on potentiometric and solubility experiments. Potentiometric titrations, utilizing ahydrogen-electrode concentration cell, were performed from 50 to 125°C at ionic strengths of 0.1, 0.3 and 1.0molal in aqueous NaCl media. Two aluminum-sulphate species, AlSO4

1and Al(SO4)22, were identified from

the titration data and the formation quotients for these species were modeled by empirical equations todescribe their temperature and ionic strength dependencies. Thermodynamic parameters for the complexationreactions were obtained by differentiating the empirical equations with respect to temperature. The thermo-dynamic quantities obtained for the formation of AlSO4

1 at 50°C and infinite dilution are: logK1 5 3.76 0.4,DH1° 5 210 6 30 kJz mol21, DS1° 5 40 6 100 Jz K21 z mol21 andDC°p 1 5 19006 800 Jz K21 z mol21;whereas the values for Al(SO4)2

2 are: logK2 5 5.6 6 0.7, DH2° 5 10 6 50 kJz mol21, DS2° 5 100 6 100J z K21 z mol21 andDC°p 2 5 28006 800 Jz K21 z mol21. A solubility study, which was undertaken to verifythe 50°C potentiometric data, was performed by reacting powdered gibbsite (Al(OH)3) with sulphate solutionsat 1023.5 and 1024 molal H1, total sulphate concentrations from 0.005 to 0.080 molal, and 0.1 and 1.0 molalionic strength in aqueous NaCl media. The results of the solubility study are in good agreement with thepotentiometric data and establish that Al-sulphate complexation substantially enhances the equilibriumsolubility of gibbsite. Copyright © 1999 Elsevier Science Ltd

1. INTRODUCTION

The thermodynamic behaviour of sulphuric acid has been thesubject of numerous studies because of the importance ofsulphur chemistry to industry ( Clegg et al., 1994; Dickson etal., 1990; Izatt et al., 1969), the well-documented environmen-tal problems associated with acid-sulphate waters ( Chapman etal., 1983; Cronan et al., 1986; Levy et al., 1992; Martin, 1994;Nordstrom, 1982; van Breemen, 1973), and the effect of sul-phate on a variety of geochemical processes ( Bloom and Erich,1987; Bradley, 1989; Cronan and Schofield, 1979; Packter andDhillon, 1969; Ridley et al., 1997). For example, field ( Alvarezet al., 1993; Chapman et al., 1983; McKnight and Bencala,1990; Monterroso et al., 1994; Nordstrom and Ball, 1986) andexperimental ( Bloom and Erich, 1987; Lydersen et al., 1991;Packter and Dhillon, 1969; Ridley et al., 1997) observationshave shown that the solubility of aluminum-bearing minerals isenhanced by the presence of sulphate in natural aqueous solu-tions. Elevated sulphate concentrations are found in a variety ofacidic surface and subsurface waters, including shallow geo-thermal systems, acid-hypersaline groundwater, acid rain andacid rock drainage ( Karlsson et al., 1988; Long et al., 1992a;Martycak et al., 1994; Nordstrom, 1982; Raymahashay, 1968;White, 1957). For example, surface and groundwaters contam-inated by acid mine drainage may have sulphate concentrationsexceeding 5100ppm with pH values,3; whereas sulphateconcentrations are typically,22ppm in near neutral pristine

waters ( Filipek et al., 1987). Moreover, in near neutral watersaluminum concentrations are low (1025 - 1028M) ( McKnightand Bencala, 1990; Nordstrom, 1982). However, aluminumconcentrations may exceed 1022.5M ( Alvarez et al., 1993;Cronan and Schofield, 1979; Cronan et al., 1990; Driscoll,1985; Lawrence et al., 1988) in acid-sulphate waters. Theoccurrence of aluminum-sulphate minerals has been widelyreported; for example, they characterize hydrothermal alter-ation zones ( Bird et al., 1989; Cunningham et al., 1984;Nordstrom, 1982; Raymahashay, 1968; Stoffregen and Alpers,1987, 1992; Zotov, 1971), occur in the supergene zone of manyore deposits ( Bird et al., 1989; Schoen et al., 1974; Stoffregenet al., 1994), and are also common products of weathering inacid-sulphate environments ( McKnight and Bencala, 1990;Monterroso et al., 1994; Nordstrom, 1982; Nordstrom et al.,1984; Van Breemen, 1973). Aluminum-sulphate minerals havealso been identified as pore-fillings and layers in near-surfaceplaya sediments of acid-hypersaline lakes ( Alpers et al., 1992;Long et al., 1992a,b; Stoffregen et al., 1994), and as a diage-netic component of sediments ( Bird et al., 1989; Cunninghamet al., 1984; Goldbery, 1980; Rouchy and Pierre, 1987). Fur-thermore, Ridley et al. (1997) showed that sulphate enhancesthe dissolution rate of gibbsite in low temperature acid waters,independently of pH and ionic strength effects.

Clearly, the geochemical behaviour of aluminum in aqueoussystems is modified by interaction with sulphate; however,there have been relatively few experimental studies to measurethe association quotients for Al-sulphate complexes at temper-atures other than 25°C, or at geochemically relevant conditions(Table 1). Moreover, there is considerable variation betweenthe experimental studies performed at 25°C: association con-stants for the formation of the AlSO4

1 species range from 101.90

*Author to whom correspondence should be addressed (E-mail:[email protected]).Present address:Department of Geosciences, Texas Tech University,P.O. Box 41053, Lubbock, TX 79409-1053, USA

Pergamon

Geochimica et Cosmochimica Acta, Vol. 63, No. 3/4, pp. 459–472, 1999Copyright © 1999 Elsevier Science LtdPrinted in the USA. All rights reserved

0016-7037/21801 $20.001 .00

459

to 103.73 ( Sharma and Prasad, 1970; Nishide and Tsuchiya,1969, respectively), and there is an even greater variation information constants for the Al(SO4)2

2 species (Table 1). Thisvariation may result partially from the large differences amongthe bisulphate dissociation constant values used to interpret thecomplexation data ( Dickson et al., 1990; Lo et al., 1982; Lutset al., 1994; Matsushima and Okuwaki, 1988). The potentio-metric study of Matsushima et al. (1988) is the only other studyof AlSO4

1 formation constants at temperatures above 70°C andas will be shown, is in good agreement with our results.

The potentiometric experiments presented in this paper wereinitiated to determine the speciation for Al-sulphate complex-ation and to examine the effects of temperature and ionicstrength on the association quotients of the Al-sulphate species.Experiments performed at 25°C were excluded from the resultsof this study, because the rather weak Al-sulphate interactioncoupled with uncertainty in the bisulphate dissociation constantresulted in an unacceptable level of accumulated error ( Dick-son et al., 1990; Lo et al., 1982; Luts et al., 1994; Matsushimaand Okuwaki, 1988). At higher temperatures the bisulphatedissociation quotients are well known as HSO4

2 is a weakeracid ( Dickson et al., 1990); therefore, extrapolation of the highertemperature Al-sulphate association quotients may give more re-liable values at 25°C than direct experimental measurement.

Solubility experiments performed at 50°C were undertakento validate the potentiometric data and to examine the effect ofsulphate on the solubility of gibbsite (and by analogy, otherAl-bearing minerals). Similar experiments could not be per-formed successfully at 25°C, as the dissolution kinetics ofgibbsite are too slow at this temperature. Finally, the thermo-

dynamic data and empirical equations presented in this paperprovide an accurate model for the speciation of Al-sulphateover a range of temperatures and ionic strengths relevant toenvironmental and geochemical conditions.

2. EXPERIMENTAL

2.1. Potentiometric Titrations

2.1.1. Materials

Three solutions at each ionic strength (0.1, 0.3 and 1.0 molal)were prepared by diluting concentrated stock solutions withdistilled, deionized water. The stock solutions of HCl, NaOH,NaCl and AlCl3 were prepared, standardized and stored fol-lowing the procedures described by Palmer and Wesolowski(1993). A stock solution of sulphuric acid (1N Baker analyzedreagent, lot H13528) was standardized by acidimetric titration(by weight) against the NaOH stock solution, and used toprepare the 0.1 and 0.3 molal ionic strength titrant solutions.The titrant solution at 1.0 molal ionic strength was preparedfrom crystalline solid Na2SO4 (Aldrich lot # TW 08314 PW),which had been dried overnight in a vacuum oven at 50°C. Alltitrant solutions were stored under a positive pressure of argon.

The initial pH of the test solutions was less than 2.5, whichincreased during the course of a titration but never exceeded 3;this range of pH minimized the formation of hydrolyzed Al-species. (Note that throughout this study pHm [ 2log[H1], inmolal concentrations ( Mesmer and Holmes, 1992)). The com-position of the test, reference and titrant solutions used in thetitrations are given in Table 2.

Table 1. Summary of the available data for the formation of Al-sulphate complexes.

log QaDH

kJ z mol21DS

J z mol21 z K21Temp

°C I. Str./Medium Reference

Al31 1 SO422 5 AlSO4

1:3.01 9.59 89.6 25 0.0 Izatt et al., 19690.437 29.01 25 1.0 M NaClO4 Lo et al., 19820.603 28.68 35 1.0 M NaClO40.831 28.39 50 1.0 M NaClO41.104 28.39 70 1.0 M NaClO43.35 6.61 86.2 25 0.0 Matsushima et al., 19883.59 30.3 162 50 0.04.08 50.9 224 75 0.04.66 68.6 273 100 0.05.34 83.3 312 125 0.03.73 25 0.0 Nishide and Tsuchiya, 19691.600 5 0.0 Sharma and Prasad, 19701.749 15 0.01.898 25 0.02.078 35 0.02.57 25 0.0 Stryker and Matijevic, 1969

Al31 1 2SO422 5 Al(SO4)2

2:4.90 12.85 136.9 25 0.0 Izatt et al., 19690.728 39.31 25 1.0 M NaClO4 Lo et al., 19820.971 45.38 35 1.0 M NaClO41.376 56.65 50 1.0 M NaClO42.006 75.70 70 1.0 M NaClO42.249 15 0.0 Sharma and Prasad, 19702.697 25 0.03.125 35 0.0

a All log Q values are expressed in terms of molality; values reported in terms of molarity were converted tomolality (Helper, 1981).

460 M. K. Ridley et al.

2.1.2. Equipment and procedure

The hydrogen-electrode, concentration cell and auxiliaryequipment used, and the titration procedures followed in per-forming the potentiometric titrations, have been described pre-viously (Giordano and Drummond, 1991; Kettler et al., 1991;Mesmer et al., 1991; Palmer and Bell, 1994). In this study, theconfiguration of the concentration cell prior to the addition ofthe sulphate titrant was:

Pt, H2 u HCl, AlCl3, NaCl i HCl, NaCl u H2, Pt

Test Reference

with the sulphate titrant being added to the Al-test solution.This cell configuration was used in order to vary the ligandconcentration widely, thereby facilitating the quantitative de-termination of the complex stoichiometries and formation con-stants. The cell potential was considered stable when threeconsecutive readings agreed to within6 0.01 mV (ca. 6 min)after an addition of titrant. All titrations performed at temper-atures above 125°C exhibited drifting cell potentials, whichwere ascribed to the formation of polynuclear Al-hydrolysisspecies, and / or precipitation of boehmite (AlOOH), therebyconstraining the upper temperature limit of this study.

2.2. Solubility Experiments

2.2.1. Materials

A set of twenty-four solutions were prepared from the con-centrated H2SO4, NaOH and NaCl stock solutions used for thepotentiometric titrations. The experimental solutions were pre-pared at ionic strengths of 0.1 and 1.0 molal, and two hydrogenion molalities of 1023.5 and 1024. At 0.1 molal ionic strength,the stoichiometric sulphate concentration was varied from0.005 to 0.030 molal, whereas at an ionic strength of 1.0 molalthe stoichiometric sulphate concentration ranged between 0.03and 0.08 molal (Table 3).

The gibbsite used in these experiments was taken from thesame batch of pure, synthetic gibbsite (Alcoa compositionC-31) as that used in earlier studies performed in this laboratory(Ridley et al., 1997; Wesolowski, 1992; Wesolowski andPalmer, 1994). This material was pretreated following theprocedures suggested by Bloom and Weaver (1982) to removeany fines and high-energy surfaces. Characterization of thegibbsite starting material was described in detail by We-solowski (1992).

2.2.2. Experimental procedure and analytical method

The experiments were performed in 50 mL, disposable,sterile, polypropylene/polyethylene syringes, which wereloaded with approximately 3 g of gibbsite and a 35–40 galiquot of experimental solution. The syringes were thenmounted on a rotating rack (6 revolutions per hour) in athermostated water bath, at 50°C6 0.05°C, and allowed toreact overnight. The initial aliquot of experimental solution wasdiscarded and replaced with fresh starting solution, then thesyringes were returned to the rotating rack in the water bath at50°C. The syringes were sampled after 25, 46, 75 and 112 days.At each sampling the outer surface of each syringe was dried,then fitted with a 0.2mm PVDF membrane filter through whicha small volume of solution was initially dispensed and dis-carded. Finally a 1 mLsample was dispensed directly into a 5mL disposable, polypropylene/polyethylene syringe attached tothe downstream end of the filter; these filtered samples showed

Table 2. Summary of the starting molal (molz kg21) solution compositions, for the potentiometric titrations of Al31 with sulphate.

Reference Test Titrant

mHCl mNaCl mHCl mAlCl3 mNaCl m(H2SO4 mNaOH mNaCl

0.005000 0.9950 0.009997 0.005002 0.9603 0.05000 — 0.99860.002500 0.0975 0.004995 0.002001 0.0830 0.02000 0.0400 0.05000.002496 0.2967 0.004951 0.003049 0.2753 0.03000 0.0600 0.30000.002496 0.2967 0.005024 0.001996 0.2822 0.03000 0.0600 0.3000

Table 3. Summary of the starting molal (molzkg21) solution compo-sitions for the dissolution of gibbsite in aqueous H2SO4-NaCl media at50°C.

Sample # mH2SO4 mNaOH mNaCl pH meas.

I 5 0.1 mS1 0.00498 0.00788 0.0854 3.478S2 0.00999 0.0169 0.0716 3.473S3 0.0150 0.0256 0.0583 3.460S4 0.0200 0.0344 0.0451 3.464S5 0.0250 0.0429 0.0318 3.475S6 0.0300 0.0513 0.0193 3.478S7 0.00501 0.00981 0.0848 4.035a

S8 0.0100 0.0197 0.0704 3.998S9 0.0150 0.0297 0.0551 4.010S10 0.0200 0.0396 0.0403 4.042S11 0.0250 0.0495 0.0259 4.000S12 0.0300 0.0593 0.0109 3.960

I 5 1.0 mS13 0.0300 0.0551 0.911 3.550S14 0.0400 0.0740 0.882 3.542S15 0.0500 0.0927 0.855 3.530S16 0.0600 0.111 0.827 3.520S17 0.0700 0.129 0.800 3.515S18 0.0797 0.147 0.769 3.507S19 0.0300 0.0597 0.910 4.102S20 0.0400 0.0796 0.880 4.148S21 0.0500 0.0996 0.850 4.135S22 0.0600 0.120 0.820 4.094S23 0.0700 0.139 0.791 4.036S24 0.0799 0.160 0.761 4.334

a Value measured after 75 days, all other values were measured after112 days.

461Aluminum-sulphate complexation in NaCl media

no visible evidence of entrained particulates or precipitates. Analiquot of each sample was immediately diluted by weight-addition with 0.01M HCl for analysis of the total dissolved Alcontent, which was measured by ion chromatography. The ionchromatographic analytical procedures used in this study arediscussed in detail by Wesolowski and Palmer (1994).

Although it was possible to calculate the initial pHm of theexperimental solutions from the known solution stoichiometry,pHm was also measured at 50°C during the course of the studyand on completion of the study, following the procedure de-scribed by Wesolowski and Palmer (1994). When collectingsamples for Al analysis, an additional filtered sample wascollected and placed in a second thermostated water bath at50°C. A Ross combination glass pH electrode was equilibratedat 50°C, then standardized with four solutions at each ionicstrength. The standard solutions at 1.0 molal ionic strengthcontained 53 1025 to 0.05 molal HCl in NaCl media, whereasthe 0.1 molal standard solutions contained 53 1025 to 1023

molal HCl in NaCl media. A calibration curve was establishedat each ionic strength by recording the potential readings (inmillivolts) of each standard solution. The potential of eachsample was then measured at 50°C and the H1 concentrationcalculated from the calibration curve at the corresponding ionicstrength. The calibration curves were checked frequently fordrift during the period required to measure all samples.

3. RESULTS

3.1. Potentiometric Titrations

Hydrogen and Al31 ions compete for association with thesulphate anions. Given the known bisulphate dissociation con-stants ( Dickson et al., 1990) and the stoichiometric concentra-tions of H1, SO4

22, and Al31, the measured molal concentra-tion of hydrogen ions in the test solutions ([Htest

1 ]) provides anaccurate measure of the association between Al31 and sulphate.The molal concentration of hydrogen ions in the test solutioncan be related to the cell potential through the Nernst equation:

2log @Htest1 # 5

2.303F

RT~E 1 ELJ! 2 log @Href

1 # (1)

where F is the Faraday constant; R is the universal gas constant;T is temperature in Kelvin; [H1ref] is the known stoichiometrichydrogen ion concentration of the reference solution; and E andELJ are the cell potential and calculated liquid junction poten-tial, respectively ( Ridley, 1997). Implicit in these calculationsis the assumption that the activity coefficient of the H1 ion isidentical in the test and reference compartments because theionic strengths are similar and the supporting electrolyte (NaCl)is in excess of the other ions. Any ion pairing of H1 and Na1

with Cl2 and SO422 is implicitly incorporated into the activity

coefficient ratio of the bisulphate dissociation reaction, whichwas studied using the same experimental and data reductionschemes ( Dickson et al., 1990).

The liquid junction potential was calculated using the fullHenderson equation (Eqn. [2–12] in Baes and Mesmer, 1976).Limiting molar conductance data are known for H1, Na1, Cl2,HSO4

2 and SO422 ( Quist and Marshall, 1965), but data for Al31

and the Al-sulphate species have not been measured over thetemperature ranges considered in this study. The limiting molar

conductance of the AlSO41 and Al(SO4)2

2 species were mod-eled using the conductance values of Na1 and Cl2, respec-tively; whereas Al31 was modeled as La31 ( Robinson andStokes, 1959). These and similar assumptions have been usedpreviously in this laboratory ( Giordano and Drummond, 1991;Kettler et al., 1991, 1992; Palmer and Bell, 1994; Palmer andDrummond, 1988; Palmer and Wesolowski, 1993), and con-tribute little to the uncertainty assigned to the liquid junctionpotential. The calculated liquid junction potentials of the titra-tions performed at 0.1m ionic strength were less than 1.8mV,whereas liquid junction potentials of less than 0.5mV weretypical for the titrations at 0.3 and 1.0m ionic strength. A liquidjunction potential of 1.8mV would contribute an uncertainty of,0.010 to the pHm of the test solution, assuming that theHenderson equation predicts ELJ to within 25% ( Mesmer andHolmes, 1992). The higher liquid junction potentials, calcu-lated at 0.1 m ionic strength reflect the greater contribution ofH1, Al31, and sulphate ions to the total solution ionic strength,relative to the contribution of these ions at 0.3 and 1.0m ionicstrength.

The hydrogen ion ([H1test]), free sulphate ([SO422]), and free

aluminum ([Al31]) concentrations were calculated using aniterative process involving the Nernst expression (Eqn. 1), thecontribution of Al-hydrolysis species ( Wesolowski andPalmer, 1994), the dissociation quotients of bisulphate in NaClmedia ( Dickson et al., 1990), and refinement of the ELJ andionic strength terms. The convergence criterion for this routinewas that ELJ changed less than 0.0001 mV between successiveiterations. A summary of the results for each titration (E, ELJ,n# , pHm, SAl and SSO4) were presented in Ridley (1997).Following this process, the degree of complexation (n#), definedconventionally ( Baes and Mesmer, 1976) as the average num-ber of sulphate ions bound per unhydrolyzed Al ion, wascalculated from

n# 5(y@Al ~SO4!y

322y#

(Al2(@Al ~OH!q32q#

5(SO4 2 ~@SO4

22# 1 @HSO42#!

(Al 2 (@Al ~OH!q32q#

5

(SO4 2 @SO422#S1 1

@H1#

Q D(Al 2 (@Al ~OH!q

32q#(2)

assuming the formation of only mono-aluminum species and nomixed hydroxy-sulphate complexes. The numerator of Eqn. 2reflects all sulphate bound to aluminum; and the denominator isthe total stoichiometric aluminum concentration,SAl, correctedfor the presence of Al-hydrolysis species. The concentration ofassociated sulphate is given by the initial stoichiometric con-centration of sulphate,SSO4, minus the calculated concentra-tion of sulphate and bisulphate; [H1] and Q represent thecalculated molality of hydrogen ions and the molal dissociationquotients of bisulphate ( Dickson et al., 1990), respectively.

The correction applied to the denominator of Eqn. 2 wasneeded in order to preserve the conventional definition of n#(Baes and Mesmer, 1976) as the average number of ligandsbound perfree metal in the solution. By this convention n#would then approach unity with increasing free sulphate con-


centration if the only significant species were AlSO41 and

Al(OH)q32q. The concentrations of Al-hydrolysis species,

S[Al(OH)q32q], were computed iteratively using the Al-hydro-

lysis formation quotients of Wesolowski and Palmer (1994).The contribution of Al-hydrolysis species to the total stoichi-ometric concentration of Al was most significant at low valuesof ionic strength and high temperatures. The concentration ofAl-hydrolysis species typically decreased during the course ofa titration. With the exception of two titrations, Al(OH)q

32q

comprised less than 2% of the total Al concentration. At 0.1and 0.3 molal ionic strength and 125°C,S[Al(OH)q

32q] con-tributed 14.04% and 9.33% to the total Al concentration, re-spectively.

The maximum n# values (n#max) calculated for each titrationvaried as a function of temperature; such that nmax values at100 and 125°C were typically.1.0, whereas at 50 and 75°Cnmaxwas typically,0.5 (Table 4). The n¯ values were regressedusing a non-linear least-squares fitting routine ( Busing andLevy, 1962) to establish the appropriate speciation and forma-tion quotients of complexes formed between Al31 and sulphateaccording to Eqn. 3.

n# (calc) 5(yQy[Al 31][SO4

22]y

(Al 2 ([Al(OH)q32q]

(3)

Equation 3 is derived from Eqn. 2 by substituting[Al(SO4)y

322y] for the formation quotient, Qy, which is definedas

Qy 5[Al(SO4)y

322y]

[Al 31][SO422]y 5 Ky

gAl 31gSO422

y

gAl(SO4)y322y

(4)

where Ky is the corresponding equilibrium constant at infinitedilution and gi are the individual ionic activity coefficients(Giordano and Drummond, 1991; Palmer and Bell, 1994;Palmer and Drummond, 1988). All n# values for the titrationsperformed at 50 to 100°C were included in the fitting routine;whereas at 125°C only n# values less than 1.6 were included inthe regression. Two mononuclear Al-sulphate complexes,AlSO4

1 and Al(SO4)22, were identified and are described by the

reaction

Al31 1 ySO422º Al(SO4)y

322y (5)

where y5 1 or 2. The n# values greater than 1.6 suggest that a1:3 species, Al(SO4)3

32, may have formed at 125°C. However,there were insufficient n# values above 1.6 to determine accurateassociation quotients for an Al(SO4)3

32 species. The molalformation quotients resulting from the least-squares fitting pro-cedure are listed for each titration in Table 4 and the agreementbetween representative experimental n# values (Eqn. 2) and themodeled n# values (Eqn. 3) is shown in Fig. 1.

3.2. Solubility Experiments

The total Al concentration at each sampling and the mea-sured pHm values are shown in Table 5, and plotted as afunction of time in Fig. 2. As apparent from Fig. 2, the Alconcentration in sulphate solutions neared equilibrium with thegibbsite within a few weeks, and the concentration of totaldissolved Al was nearly constant from the second sampling (46days) until the study was terminated (112 days). It must benoted that all experiments were equilibrated from undersatura-tion (SAl 5 zero), and that this study was not performed toexamine the kinetics of gibbsite dissolution, but to verify thepotentiometric data for the formation of Al-sulphate complexesat 50°C. However, Fig. 2 suggests that the higher ionic strengthsolutions (I5 1.0 m) were slower to equilibrate.

In the absence of sulphate as a complexing ligand, thedissolution of gibbsite is controlled by

Al(OH)3 1 3H1º Al31 1 3H2O (6)

and the hydrolysis of Al31

Al31 1 qH2Oº Al(OH)q32q 1 qH1 (7)

The concentration of total dissolved Al would, therefore,equal {[Al31] 1 S[Al(OH)q

32q]}, in an aqueous solution inequilibrium with gibbsite. In this study, the only significantAl-hydrolysis species present was AlOH21 and all other spe-cies were negligible between pHm values of 3 and 4.5 ( We-solowski and Palmer, 1994). The sum of [Al31] and [AlOH21]calculated at pHm 3.5 and 4, using the equilibrium quotients ofWesolowski and Palmer (1994) for Eqns. 6 and 7, are 0.2 to 1.0log units lower than the concentrations of dissolved Al mea-sured in the present study. Clearly, the solubility of gibbsite issignificantly enhanced by the presence of sulphate. Further-more, it is apparent from Fig. 2 that the concentration of totaldissolved Al is not only a function of pHm, but also a functionof the stoichiometric concentration of sulphate. The differencebetween the measured concentrations of Al and the computed

Table 4. Experimental association quotients for the formation ofAl(SO4)y

3–2y complexes.a

log Q1 log Q2

Temp.°C

Imol z kg21 ELJ max n#max

2.51 6 0.03 3.866 0.22 49.9 0.099 1.28 0.682.30 6 0.02 49.8 0.100 21.26 0.462.35 6 0.02 74.8 0.099 21.08 0.532.18 6 0.11 5.236 0.04 99.7 0.098 20.90 0.792.58 6 0.04 5.316 0.03 99.5 0.098 20.95 0.863.98 6 0.03 6.906 0.03 124.3 0.098 21.15 1.894.18 6 0.06 7.006 0.08 124.5 0.098 21.17 1.681.88 6 0.01 49.9 0.307 20.42 0.321.67 6 0.02 50.0 0.306 20.42 0.211.57 6 0.04 50.0 0.307 20.42 0.251.40 6 0.07 3.436 0.10 74.8 0.309 20.35 0.211.94 6 0.01 74.9 0.312 20.37 0.371.66 6 0.07 4.686 0.02 99.7 0.306 20.30 1.072.22 6 0.03 4.366 0.05 99.9 0.311 20.31 1.083.56 6 0.06 6.296 0.05 124.5 0.304 20.32 2.023.18 6 0.03 6.096 0.02 125.0 0.301 20.29 2.141.36 6 0.05 2.366 0.31 50.0 1.026 20.25 0.301.28 6 0.04 2.846 0.07 50.0 1.016 20.25 0.371.32 6 0.04 2.166 0.43 74.8 1.016 20.22 0.321.73 6 0.04 2.526 0.58 99.7 1.024 20.18 0.421.28 6 0.06 3.306 0.04 99.7 1.021 20.18 0.662.03 6 0.03 4.296 0.03 124.7 1.007 20.16 1.63

a The error estimates associated with Qy result from the non-linearleast-squares fitting routine, and do not represent titration reproducibility.


values of {[Al31] 1 [AlOH21]} should, therefore, correspondto the Al-sulphate species (S[Al(SO4)y

322y]) identified in thepotentiometric study. The total dissolved Al in the sulphatesolutions can then be described by the mass balance equation:

(Al 5 [Al 31] 1 ([Al(OH)q32q] 1 ([Al(SO4)y

322y] (8)

4. DISCUSSION

4.1. Potentiometric Results

To provide empirical equations that describe the temperatureand ionic strength dependencies of the formation quotients foreach complex (Table 4) obtained from the potentiometric titra-tions, the non-linear least-squares fitting routine of Busing andLevy (1962) was used. Criteria used to determine the empiricalequations are outlined by Giordano and Drummond (1991), andPalmer and Bell (1994). The empirical equations that bestdescribe the AlSO4

1 and Al(SO4)22 association quotients and

require the fewest number of adjustable parameters are definedby:

log Qy 52Dz2Aw

ln(10) H ÎI

(1 1 bÎI)1 Sa

bD ln(1 1 bÎI)J 1 p1

1p2

T1 p3T 1 p4I

2 1 p5I 1 p6I/T (9)

where the first term is the extended Debye-Hu¨ckel expressionfrom Pitzer (1973), and includesa (2.0), b (1.2) and Aw takenfrom Bradley and Pitzer (1979);Dz2 has values of212 and216 for the first and second association quotients, respectively;I is the stoichiometric molal ionic strength; and T is the abso-lute temperature. Values for the variable parameters (p126) in

Fig. 1. The agreement between the n# values determined from Eqn. 2 (closed symbols) and calculated n# values from theleast-squares fitting routine (open symbols and solid lines), when considering AlSO4

1 and Al(SO4)22 complexes, are shown

for titrations at 0.1 molal ionic strength. The n# values are shown as symbols at temperatures of (F) 50; (�) 75; (■) 100;and (r) 125°C.

Table 5. Gibbsite solubility in aqueous H2SO4-NaCl media at 50°C.

Smp.#

2log[H1]meas. log[(Almeas.]

Day 75 Day 112 Day 25 Day 46 Day 75 Day 112

I 5 0.1 mS1 3.480 3.478 23.250 23.250 23.261 23.274S2 3.480 3.473 23.084 23.074 23.077 23.102S3 3.463 3.460 22.917 22.885 22.891 22.943S4 3.476 3.464 22.818 22.802 22.807 22.794S5 3.476 3.475 22.714 22.649 22.688 22.697S6 3.464 3.478 22.625 22.580 22.596 22.608S7 4.035 NS 24.807 24.775 24.818 NSS8 3.966 3.998 24.506 24.468 24.479 24.543S9 3.996 4.010 24.401 24.350 24.372 24.401S10 4.030 4.042 24.364 24.328 24.324 24.380S11 3.994 4.000 24.182 24.139 24.119 24.146S12 3.948 3.960 23.947 23.885 23.906 23.914

I 5 1.0 mS13 3.512 3.550 22.947 22.859 22.857 22.878S14 3.500 3.542 22.833 22.754 22.776 22.775S15 3.484 3.530 22.754 22.676 22.685 22.684S16 3.463 3.520 22.733 22.588 22.603 22.593S17 3.454 3.515 22.602 22.519 22.541 22.534S18 3.444 3.507 22.541 22.458 22.453 22.452S19 4.023 4.102 24.572 24.301 24.278 24.307S20 4.067 4.148 24.513 24.383 24.389 24.415S21 4.036 4.135 24.394 24.271 24.270 24.300S22 4.001 4.094 24.200 24.093 24.097 24.116S23 3.957 4.036 24.006 23.922 23.932 23.939S24 4.217 4.334 ND 24.569 24.564 24.732

NS 5 No sample; ND5 Not detected.


Eqn. 9 are given in Table 6. The terms p1 – p3 of Eqn.9 definethe equilibrium constants (log Ky), with terms p4 – p6 beingfunctions of ionic strength required to fit the deviation from theDebye-Huckel expression at finite ionic strengths. The good-ness of fit between the association quotients listed in Table 4and Eqn. 9 are shown for both species in Fig. 3. Agreement

between the Q1 values and the modeled curves is adequate (Fig.3). However, reproducibility between duplicated titrations ispoorer than expected (Table 4), particularly when comparingthe data obtained in this study with similar recent measure-ments for the complexation of Al31 with malonic acid ( Ridleyet al., 1998). This is attributed to the fact that HSO4

2 is arelatively strong acid at low temperatures and elevated ionicstrengths, and computed Qy values are strongly influenced bysmall uncertainties in temperature and solution composition.The fit for the di-sulphate species could not be constrained asrigorously as was possible for Q1, because of the limitednumber of Q2 values at 50 and 75°C and their large associatederrors (Table 4). The difficulty in determining Q2 at 50 and

Fig. 2. Plots of the total dissolved aluminum, representing the dissolution of gibbsite as a function of time, in H2SO4-NaClmedia at 50°C. The solutions in plots (a) and (b) had an ionic strength of 0.1m and pHm of approximately 3.5 and 4,respectively (solutions S1 to S12, in Tables 3 and 5); whereas the ionic strength of the solutions represented by plots (c)and (d) was 1.0 m with pHm of approximately 3.5 and 4, respectively (solutions S13 to S24, in Table 3 and 5). The symbolsin plots (a) and (b) represent total sulphate concentrations of (F) 0.005 m; (E) 0.01m; (■) 0.015 m; (h) 0.02 m; (Œ) 0.025m; and (‚) 0.03 m; whereas in plots (c) and (d) the symbols define total sulphate concentrations of (F) 0.03 m; (E) 0.04m; (■) 0.05 m; (h) 0.06 m; (Œ) 0.07 m; and (‚) 0.08 m.

Table 6. Parameters for equation 9.

Qn Dz2 p1 p2 p3 p4 p5 p6

1 212 295.7623 16330.1 0.15146121.18203 2.47143 —2 216 2139.898 23211.4 0.22789222.89218 — 1669.08


75°C may be ascribed to the low n#max values obtained at thesetemperatures, consistent with the di-sulphate complex being aminor species. Association quotients for both Al-sulphate spe-cies, calculated from Eqn. 9, are given in Tables 7 and 8 forselect temperatures and ionic strengths. At all conditions pre-sented in Tables 7 and 8, the stepwise association quotients forAl(SO4)2

2 formation (i.e. AlSO41 1 SO4

22 ^ Al(SO4)22) are

less than (within error) the equilibrium quotients for the for-mation of AlSO4

1; this was not the case for all experimentalassociation quotients presented in Table 4.

Differentiation of Eqn. 9 with respect to temperatureyields the thermodynamic quantitiesDH, DS andDCp; val-ues for these parameters are given in Tables 7 and 8. Theenthalpy, entropy and heat capacity values are generallylarge and positive, and increase with increasing temperature.

However, large errors are associated with these thermody-namic values arising from the large uncertainties in theequilibrium quotients (Table 4). Despite these large errors, itis apparent that the change inDS to more positive values isresponsible for the complexation reactions, which outweighsthe positive change inDH values.

4.2. Solubility Results

Two approaches were followed utilizing the 50°C gibbsitesolubility data (Table 5) to validate the Al-sulphate formationquotients obtained from the potentiometric titrations (Tables 7and 8). In the first approach, the concentration of total dissolvedAl in equilibrium with gibbsite was calculated for each solutionin the 50°C solubility experiment using an iterative process,and a stepwise refinement of ionic strength. This routine in-cluded the Al-sulphate association quotients determined in thisstudy (Tables 7 and 8), the equilibrium quotients of We-solowski and Palmer (1994) for the solubility of gibbsite andthe formation of Al-hydrolysis species in NaCl media, thestoichiometric concentration of sulphate (Table 3), and themeasured pHm values (Table 5). The results of this routine aresummarized in Table 9.

The deviation between the calculated Al concentrations andthe measured Al concentrations (logSAlmeas2 logSAlcalc) areshown for all solubility experiments in Fig. 4. With the excep-tion of one experiment, the maximum deviation is less than60.2 log units; which is well within the error range of theAl-sulphate formation quotients, log Q1 and log Q2, at 50°C

Fig. 3. The association quotients, log Qy, for the formation ofAl(SO4)y

322y species are plotted as a function of temperature (°C),where the symbols define the measured log Qy values at ionic strengthsof (F) 0.1, (■) 0.3, and (Œ) 1.0 molal. The curves were computed fromEqn. 9, and the dashed lines represent association constants at infinitedilution, also computed from Eqn. 9. Also plotted are the AlSO4

1

association constants at infinite dilution ({) reported by Matsushima etal. (1988).

Table 7. Summary of the thermodynamic quantities for the formationof AlSO4

1 at the saturation vapour pressure of water.a

T°C log Q1

DHkJ z mol21

DSJ z K21 z mol21

DCp

J z K21 z mol21

I 5 0.050 3.76 0.4 2106 30 406 100 19006 80075 3.96 0.4 406 10 2006 100 20006 900

100 4.56 0.4 906 20 3006 100 22006 1000125 5.66 0.4 1506 40 5006 100 23006 1000

I 5 0.150 2.36 0.2 2106 30 106 100 19006 80075 2.46 0.2 406 10 2006 100 20006 900

100 2.96 0.2 906 20 3006 100 22006 1000125 3.86 0.4 1506 40 4006 100 23006 1000

I 5 0.250 2.06 0.2 2106 30 86 100 19006 80075 2.06 0.2 406 10 2006 100 20006 900

100 2.66 0.2 906 20 3006 100 22006 1000120 3.46 0.3 1506 40 4006 100 23006 1000

I 5 0.550 1.66 0.3 2106 30 16 100 19006 80075 1.66 0.3 406 10 1006 100 20006 900

100 2.16 0.3 906 20 3006 100 22006 1000125 2.96 0.4 1506 40 4006 100 23006 1000

I 5 1.050 1.26 0.4 2106 30 27 6 100 19006 80075 1.26 0.4 406 10 1006 100 20006 900

100 1.66 0.4 906 20 3006 100 22006 1000125 2.36 0.4 1506 40 4006 100 23006 1000

a The errors listed represent three times the standard deviation.


(Tables 7 and 8). There is a slight systematic trend between theresiduals and the total sulphate concentration at 1.0m ionicstrength (Table 9), suggesting that the Al-sulphate formationquotients obtained from the potentiometric study are slightlyweaker than the solubility measurements would imply. How-ever, this trend is small and well within the error range of theformation quotients.

In the second approach, equilibrium quotients for Al-sul-phate complexes were determined directly from the gibbsitesolubility data. This involved determining the ‘excess’ dis-solved Al (Eqn. 8) by considering the formation of Al-sulphatecomplexes, including AlSO4

1 and Al(SO4)22 species. The sol-

ubility data from each experimental solution were initiallymodeled assuming the formation of a single complex: either amonosulphate (log Q1) or disulphate (log Q2) species. In addi-tion, the solubility data were fitted with both monosulphate anddisulphate species; in this case the formation quotients weredetermined by iteration. Results of the three fitting routines arepresented in Table 10. For each Al-sulphate species considered,mean values for log Qy were calculated from the equilibriumquotients (Table 10). These mean Al-sulphate equilibrium quo-tients, and equilibrium quotients for the solubility of gibbsiteand the formation of Al-hydrolysis species ( Wesolowksi andPalmer, 1994), were then used to calculate gibbsite solubilitiesfor each experimental solution. The differences between calcu-lated Al concentrations and the total measured Al concentra-tions are presented in Table 10 and shown as deviation plots inFig. 5.

The Al-sulphate species best able to account for the gibbsite

solubility data are indicated by low mean differences betweenthe calculated and measured Al concentrations (Table 10), andby random deviations with respect to sulphate concentrationsand pHm (Fig. 5). As can be seen from Table 10, the deviationsare typically less than 0.15 log units for all Al-sulphate spe-ciation schemes considered. However, there is a slight system-atic trend between the residuals and the total sulphate concen-tration at 1.0 molal ionic strength (Fig. 5).

The solubility data at 0.1 molal ionic strength are bestdescribed by the formation of both AlSO4

1 and Al(SO4)22

complexes. When fitting the solubility data with two Al-sul-phate complexes a mean deviation between the calculated andmeasured Al-concentrations of 0.042 was obtained, which isbetter than the fits obtained for single Al-sulphate complexes(Table 10, Fig. 5). In addition, the degree of randomness isgreater when considering both Al-sulphate species (Table 10,Fig. 5). For all fits, there is close agreement between theequilibrium quotients determined in the potentiometric studyand the solubility experiments: agreement is within the errorrange of the Al-sulphate formation quotients presented in Ta-bles 7 and 8. Equilibrium quotients obtained for the formationof AlSO4

1 and Al(SO4)22 complexes at 0.1m ionic strength

were 2.13 and 3.68 for log Q1 and log Q2, respectively; whichcompares closely with the potentiometric data at the sameconditions: log Q1 5 2.3 6 0.2, and log Q2 5 3.9 6 0.6 at50°C.

The solubility of gibbsite at 1.0 molal ionic strength was alsobest accounted for by the formation of two Al-sulphate species,AlSO4

1 and Al(SO4)22. The values of log Q1 and log Q2 were

1.52 and 2.78, respectively, which are in good agreement withthe potentiometric values of 1.26 0.4 and 2.86 0.5. Whenfitting the solubility data with both AlSO4

1 and Al(SO4)22

species a mean deviation value of 0.075 was determined, whichis similar to the fit obtained for a single AlSO4

1 species.However, in the latter fit a strong pHm trend is apparent (Fig.5). Moreover, the mean formation quotient is 0.6 log unitshigher than the value obtained by potentiometry (Tables 7 and10), suggesting a second Al-sulphate species is required toadequately fit the solubility data.

Clearly, the two approaches followed to evaluate the agree-ment between the 50°C gibbsite solubility data and results ofthe potentiometric titrations are quantitatively consistent. Thegibbsite solubility data strongly support the equilibrium quo-tient values determined by potentiometry for the formation ofAlSO4

1 and Al(SO4)22 at 50°C.

4.3. Comparison with Literature Data

The equilibrium constants for AlSO41 presented in Table 7

calculated from Eqn.9, are in good agreement with the 50 to125°C data of Matsushima et al. (1988) (see Fig. 3), whopresent the only high temperature results (Table 1). Matsus-hima et al. (1988) used a hydrogen-electrode concentration cellmodeled after the Oak Ridge design. Their study was per-formed at 1.0 m ionic strength in KCl media, but they did notprovide sufficiently detailed experimental data to enable adirect comparison between the molal formation quotients oftheir study with our results in 1.0 m NaCl media. There is amarked disagreement at 25°C, where log K5 4.16 (Eqn.9),which is approximately 0.7 log units higher than the equivalent

Table 8. Summary of the thermodynamic quantities for the formationof Al(SO4)2

2 at the saturation vapour pressure of water.a

T°C log Q2

DHkJ z mol21

DSJ z K21 z mol21

DCp

J z K21 z mol21

I 5 0.050 5.66 0.7 106 50 1006 100 28006 80075 6.16 0.4 806 30 4006 100 30006 900

100 7.36 0.3 1606 10 6006 100 33006 900125 9.16 0.2 2506 30 8006 100 35006 1000

I 5 0.150 3.96 0.6 106 50 1006 100 28006 80075 4.36 0.3 806 30 3006 100 30006 900

100 5.36 0.2 1606 10 5006 100 33006 900125 6.96 0.2 2406 30 7006 100 35006 1000

I 5 0.250 3.66 0.6 56 50 1006 100 28006 80075 3.96 0.2 806 30 3006 100 30006 900

100 4.96 0.1 1606 10 5006 100 33006 900125 6.46 0.1 246 30 7006 100 35006 1000

I 5 0.550 3.46 0.6 25 6 50 506 100 28006 80075 3.66 0.2 706 30 3006 100 30006 900

100 4.46 0.2 1506 10 5006 100 33006 900125 5.76 0.2 2306 30 7006 100 35006 1000

I 5 1.050 2.86 0.5 2206 40 2106 100 28006 80075 2.76 0.2 506 30 2006 100 30006 900

100 3.36 0.2 1306 10 4006 100 33006 900125 4.46 0.2 2206 30 6006 100 35006 1000

a The errors listed represent three times the standard deviation.


Fig. 4. The deviation between measured and calculated Al concentrations (logSAlmeas2 logSAlcalc) plotted as a functionof the contribution of Al-sulphate species to the concentration of measured Al, presented as a percentrage. The procedurefollowed to calculate Al concentrations and the concentration of Al-sulphate species is outlined in the text. The symbolsdefine solutions of (F) 0.1 m ionic strength, pHm 5 3.5; (�) 0.1 m ionic strength, pHm 5 4; (V) 1.0 m ionic strength,pHm 5 3.5; and (¹) 1.0 m ionic strength, pHm 5 4.

Table 9. The distribution of Al species in Sulphate-NaCl media at 50°C; computed using the AlSO41 and Al(SO4)2

2 formation quotients determinedin the potentiometric study (Tables 7 and 8).

Smp.# pHm

log X

Deltac Ion.Str.[Al 31] [(Al(OH)q] [AlSO41] [Al(SO4)2

2] [(Alcalc]a [(Almeas]

b

I 5 0.1 mS1 3.478 23.51 24.63 23.55 24.33 23.18 23.27 20.09 0.1001S2 3.473 23.50 24.62 23.24 23.73 22.96 23.10 20.14 0.0998S3 3.460 23.46 24.59 23.03 23.36 22.76 22.94 20.18 0.0995S4 3.464 23.47 24.60 22.93 23.13 22.64 22.79 20.15 0.0991S5 3.475 23.51 24.62 22.86 22.98 22.56 22.69 20.14 0.0984S6 3.478 23.52 24.63 22.80 22.84 22.47 22.61 20.13 0.0983S7 4.035 25.19 25.72 25.17 25.91 24.79 24.82 20.03 0.0997S8 3.998 25.07 25.65 24.76 25.20 24.47 24.54 20.07 0.1002S9 4.010 25.11 25.67 24.62 24.89 24.33 24.40 20.07 0.0998S10 4.042 25.21 25.73 24.59 24.73 24.28 24.38 20.09 0.1000S11 4.000 25.08 25.65 24.37 24.42 24.04 24.14 20.10 0.1003S12 3.960 24.96 25.57 24.17 24.14 23.82 23.91 20.10 0.1000

I 5 1.0 mS13 3.550 23.28 24.60 23.60 23.55 22.97 22.88 0.09 0.9997S14 3.542 23.26 24.59 23.45 23.29 22.84 22.77 0.07 0.9992S15 3.530 23.22 24.56 23.32 23.06 22.71 22.68 0.02 0.9995S16 3.520 23.19 24.54 23.22 22.88 22.59 22.59 0.00 0.9995S17 3.515 23.18 24.53 23.14 22.74 22.49 22.53 20.04 0.9994S18 3.507 23.15 24.52 23.06 22.61 22.39 22.45 20.06 0.9946S19 4.102 24.94 25.69 25.24 25.18 24.59 24.30 0.28 1.0001S20 4.148 25.08 25.78 25.25 25.07 24.62 24.41 0.20 0.9999S21 4.135 25.04 25.75 25.12 24.84 24.48 24.30 0.18 1.0001S22 4.094 24.91 25.67 24.91 24.56 24.26 24.11 0.15 0.9999S23 4.036 24.74 25.56 24.67 24.25 24.01 23.94 0.07 1.0002S24 4.334 25.63 26.14 25.51 25.03 24.81 24.73 0.08 1.0001

a Total Al concentration calculated using the procedure outlined in the text.b Total Al measured after 112 days (Table 5).c Delta 5 log (Almeas2 log (Alcalc.


value of Matsushima et al. (1988). This difference may indicatea problem with the extrapolation of Eqn. 9 below 50°C. Noattempt was made to extrapolate Q2 values from Eqn. 9 totemperatures below 50°C; because the number of Q2 values at50 and 75°C were limited and their associated errors were large(Table 4).

4.4. Aluminum Speciation in Aqueous Fluids

The empirical equations presented in this study enable thespeciation of Al-sulphate complexes to be modeled over a widerange of geochemical and environmental conditions. The re-sults of such computations at 50 and 100°C and 1.0 molal ionicstrength are shown in Fig. 6(a–c); the formation quotients forAl(SO4)y

322y were taken from this study, whereas all otherequilibrium quotients used to construct the plots were takenfrom Wesolowski and Palmer (1994). Conditions equivalent tothe 50°C solubility experiments are shown in Figs. 6a and 6b.Typically the dominant Al-species present in solution are theAl-sulphate complexes, even at fairly low concentrations ofsulphate. The significant effect of small changes in pHm on thesolubility of gibbsite, hence the total dissolved aluminum con-centration, is also apparent. Similarly, temperature has a con-

siderable impact on the dissolution of gibbsite (and by analogyother aluminum-bearing minerals), as shown by the speciationat 100°C (Fig. 6c). Clearly, the relative enhancement ofgibbsite dissolution is greater at 100°C than at 50°C, which

Fig. 5. The deviation between measured and calculated Al concen-trations (logSAlmeas2 logSAlcalc) plotted as a function of the contri-bution of Al-sulphate species to the concentration of measured Al,presented as a percentrage. The procedure followed to calculate Alconcentrations and the concentration of Al-sulphate species is outlinedin the text. The Al-sulphate species considered in each plot were: (a)the monosulphate (AlSO4

1) complex; (b) the disulphate (Al(SO4)22)

complex; and (c) both AlSO41 and Al(SO4)2

2 complexes. The symbolsdefine solutions of (F) 0.1m ionic strength, pHm 5 3.5; (�) 0.1m ionicstrength, pHm 5 4; (V) 1.0m ionic strength, pHm 5 3.5; and (¹) 1.0mionic strength, pHm 5 4.

Table 10. Formation quotients for aluminum sulphate complexescomputed for the 50°C solubility experiment.

Smp.#

AlSO41 a Al(SO4)2

2 b AlSO41 and Al(SO4)2

2

log Q1 Deltac log Q2 Deltac log Q1 log Q2 Deltac

I 5 0.1 mS1 2.15 20.075 4.52 0.087 2.07 3.1720.014S2 2.18 20.089 4.26 0.024 2.01 3.2920.033S3 2.20 20.090 4.11 20.053 1.94 3.28 20.055S4 2.30 20.026 4.10 20.061 2.03 3.60 20.018S5 2.37 0.024 4.08 20.076 2.08 3.69 0.005S6 2.41 0.053 4.04 20.098 2.08 3.70 0.008S7 2.32 20.006 4.62 0.141 2.26 4.17 0.052S8 2.33 20.003 4.34 0.077 2.20 3.90 0.045S9 2.41 0.056 4.24 0.027 2.24 3.92 0.075S10 2.44 0.077 4.14 20.055 2.21 3.84 0.061S11 2.47 0.110 4.08 20.111 2.19 3.81 0.058S12 2.52 0.160 4.05 20.135 2.22 3.83 0.074

mean 2.34 0.064 4.21 0.079 2.13 3.68 0.042

I 5 1.0 MS13 1.71 20.063 3.27 0.078 1.55 2.81 0.009S14 1.71 20.068 3.15 0.021 1.49 2.7020.016S15 1.70 20.083 3.04 20.047 1.38 2.56 20.055S16 1.71 20.082 2.97 20.094 1.30 2.50 20.078S17 1.70 20.091 2.90 20.147 1.15 2.41 20.111S18 1.72 20.077 2.87 20.170 1.08 2.42 20.118S19 2.02 0.135 3.54 0.261 1.94 3.37 0.201S20 1.93 0.080 3.33 0.146 1.80 3.11 0.123S21 1.93 0.087 3.24 0.089 1.76 3.02 0.102S22 1.93 0.091 3.16 0.031 1.71 2.94 0.076S23 1.87 0.043 3.03 20.074 1.54 2.77 20.002S24 1.92 0.085 3.02 20.083 1.58 2.80 0.012

mean 1.82 0.082 3.13 0.103 1.52 2.78 0.075

a Log Q1 computed for the reaction Al31 1 SO422 ^ AlSO4

1,assuming only the monosulphate complex forms.

b Log Q2 computed for the reaction Al31 1 2SO422 ^ Al(SO4)2

2,assuming only the disulphate complex forms.

c Delta 5 log (Almeas2 log (Alcalc.


results from the increase in stability of the Al-sulphate forma-tion quotients with increasing temperature. However, the con-centration of total dissolved Al decreases, because of the de-crease in gibbsite solubility as temperature increases at constant

pHm. Furthermore, changes in ionic strength will influence thesolubility of gibbsite and Al-sulphate complexation. Thus, nu-merous factors can affect the concentration of total dissolvedAl and the Al-speciation of a solution.

5. CONCLUSIONS

This study presents the results of potentiometric titrationsperformed to determine the speciation and molal formationquotients for complexes formed between Al and sulphate. TwoAl-sulphate species were identified, AlSO4

1 and Al(SO4)22,

from this study. The Al-sulphate formation quotients at 50°Cwere shown to be in good agreement with the gibbsite solubil-ity study. The formation quotients of AlSO4

1 and Al(SO4)22

were modeled as a function of temperature and ionic strength,and the thermodynamic quantitiesDH, DS andDCp were de-termined for both species. The equilibrium quotients calculatedfor the formation of AlSO4

1 were in quantitative agreementwith the 50 to 125°C data of Matsushima et al. (1988).

The gibbsite solubility study clearly indicated that the pres-ence of sulphate in an aqueous solution substantially enhancesthe dissolution of gibbsite (and by analogy, other aluminum-containing minerals), thus suggesting that the formation ofAl-sulphate complexes may play an important role in control-ling the solubility of Al-bearing minerals. In addition, theempirical equations and thermodynamic data provide an accu-rate model that can be used to speciate aluminum in a varietyof aqueous environments containing sulphate (at least at$50°C), and therefore can be used to model the effects ofsulphate on geochemical and environmental processes.

Acknowledgments—This research was sponsored by the Office of BasicEnergy Sciences, U.S. Department of Energy, under contract numberDE-AC05-960R22464 with Oak Ridge National Laboratory, managedby Lockheed Martin Energy Research Corp. and NSF grant EAR-9317075. The thorough reviews by J. B. Fein and an anonymousreviewer greatly improved the manuscript.

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association quotients of aluminum sulphate complexes in nacl media from 50 to 125°c: results of a...

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