ettringite solubility and geochemistry of the ca oh –al so –h o...

Ž .Chemical Geology 148 1998 1–19

Ettringite solubility and geochemistry of thež / ž /Ca OH –Al SO –H O system at 1 atm pressure and 298 K2 2 4 3 2

Satish C.B. Myneni a,), Samuel J. Traina a,b, Terry J. Logan a

a EnÕironmental Sciences, School of Natural Resources, 202 Kottman Hall, 2021 Coffey Road, OH, USAb Department of Geological Sciences, Orton Hall, South OÕal Mall, The Ohio State UniÕersity, Columbus, OH 43210, USA

Received 9 January 1997; accepted 26 September 1997

Abstract

Ž Ž . Ž . .The solubility and weathering reactions of ettringite, Ca Al SO OH P26H O , were used to study the geochemi-6 2 4 3 12 2Ž . Ž .cal equilibria of the Ca OH –Al SO –H O system at environmental pH conditions. Ettringite is a stable mineral above a2 2 4 3 2

Ž .pH of 10.7 and dissolved congruently with a log K of y111.6 "0.8 . Between pH 10.7 and 9.5, ettringite underwentsp

incongruent dissolution to gypsum and Al-hydroxides and controlled Ca2q, Al3q, and SO2y activities. At near neutral pH,4

Al-hydroxy sulfates precipitated in addition to gypsum and Al-hydroxide. These Al-hydroxy sulfate phases exhibitedprismatic and anhedral shapes and had variable AlrS ratios. In addition, some new poorly crystalline Ca–Al-hydroxy sulfate

Ž . 2q 3q 2yphases were identified in microscopic studies when the pH was acidic pH;5 . The activities of Ca , Al , and SO4Ž . Ž .suggest that the geochemistry of the Ca OH –Al SO –H O system in the pH range of 7 to 10 is simple and its2 2 4 3 2

Ž . Ž .component Ca OH –SO –H O and Al SO –H O systems behave independently of each other. The precipitation of2 3 2 2 4 3 2

Al-hydroxy sulfates below pH 7.0 significantly influenced Ca2q and SO2y activities. This effect was pronounced when4Ž .Ca–Al-hydroxy sulfate phases started precipitating pH-5.0 . The lack of thermodynamic data on the newly identified Al,

Ž . Ž .and Ca–Al-hydroxy sulfates makes it difficult to interpret the geochemistry of Ca OH –Al SO –H O system for2 2 4 3 2

pHF5.0. Reaction path calculations conducted using the EQ6 computer code predicted ion activities close to theexperimental values above pH 5.0. The observed differences between thermodynamic modelling and actual experimentaldata below this pH can be explained by the formation of Al–rCa–Al-hydroxy sulfate phases in the system, as detected byelectron microscopy and X-ray elemental analysis. These reactions are relevant and useful to the prediction of Al, and Cageochemistry in natural systems. q 1998 Published by Elsevier Science B.V. All rights reserved.

Keywords: Ettringite; Solubility; Al geochemistry; Al-sulfates

1. Introduction

Alkaline materials such as cements, coal combus-Ž .tion residues, flue gas desulphurization FGD and

) Corresponding author. Mail Stop 90-1116, Earth SciencesDivision, Ernest Orlando Lawrence Berkeley National Laboratory,Berkeley, CA 94720.

cement-based radioactive and other contaminant-so-lidification by-products are widespread around theworld. These materials are highly reactive whenexposed to natural waters, and are observed to mod-ify soil element dissolution patterns and to controlsurface and sub-surface water quality in their vicinityŽ .Mattigod et al., 1990; Fowler et al., 1993 . Weather-

Ž .ing of these materials produces high pH )10 and

0009-2541r98r$19.00 q 1998 Published by Elsevier Science B.V. All rights reserved.Ž .PII: S0009-2541 97 00128-9

( )S.C.B. Myneni et al.rChemical Geology 148 1998 1–192

dissolved Ca2q, Al3q, and SO2y ion concentrations4Ž Ž . Ž .and precipitation of ettringite Ca Al SO OH6 2 4 3 12

.P26H O as one of the dominant secondary mineral2Žphases Mattigod et al., 1990; Damidot et al., 1992;

Damidot and Glasser, 1993; Fowler et al., 1993;.Myneni, 1995 . Researchers have also shown that

ettringite formation is concomitant with the reduc-Ž .tion of leachate trace elements Fowler et al., 1993 .

In addition, the natural occurrences of ettringite arealso reported from alkaline evaporites in JordanŽ .Linklater et al., 1996 .

Currently, studies on calcareous alkalinematerial-weathering and constituent trace elementfate are limited by the wide range of literature valueson ettringite solubility. Moreover, no information isavailable on ettringite weathering around neutral pH.Estimation of ettringite solubility in the alkaline andneutral pH range is important to the interpretation ofthe geochemistry and mineral equilibria in the

Ž . Ž . ŽCa OH –Al SO –H O ternary system Hills et2 2 4 3 2.al., 1993; Myneni, 1995 . Knowledge of these reac-

tions is also important in systems such as mixed-waste disposal sites where these alkaline materialsare assimilated with acidic mine spoils to alleviate

Ž .acid generation Fowler et al., 1993 . To better un-derstand ettringite geochemistry, the objectives of

Ž .this research are focused at the: 1 estimation of thestandard state free energy value for ettringite from

Ž .solubility experiments; 2 evaluation of the ettrin-Ž .gite solubility at near neutral pH; and 3 developing

Ž . Ž .a geochemical model for the Ca OH –Al SO –2 2 4 3

H O system for natural environmental conditions.2

1.1. Solubility and standard state free energy offormation of ettringite

Congruent dissolution of ettringite can be de-scribed by the equation:

Ca Al SO OH P26H Om6Ca2qq2Al3qŽ . Ž . 126 2 4 23

q3SO2y q12OHyq26H O 1Ž .4 2

This reaction can also be written by substituting3q Ž Ž .y. 3qhydroxylated Al species e.g., Al OH for Al4

in the products. Although this results in differentsolubility constant for ettringite dissolution, they are

Ž .related to Eq. 1 through Al-hydrolysis constants,and thus there is no effect on the calculated standardstate free energy of formation.

According to the law of mass action, the equilib-Ž .rium constant for Eq. 1 is given by:

6 2 3 12 262q 3q 2y yCa Al SO OH H OŽ . Ž . Ž . Ž .Ž .4 2K s1 EttringiteŽ .

Ž . Ž .or when ettringite s H O s12

6 2 3 12ett . 2q 3q 2y yK s K s Ca Al SO OHŽ . Ž . Ž .Ž .1 sp 4

where ett.K is the solubility product constant forsp

ettringite dissolution, and the parentheses denote ac-tivities. The equilibrium constant is related to the

Ž .free energy change of the reaction DG8 . Accordingto Hess’ Law:

DG8syRT lnett .K sDG 8yDG 8sp p r

where R is the universal gas constant, T the absolutetemperature and DG 8 and DG 8 are the respectivep r

standard state Gibbs free energies of formation ofproducts and reactants. In other words,

yRT lnett .K s 6DG 2q8q2DG 3q8q3DG 2y8Žsp Ca Al SO4

q12DG y8 q26DG 8 yDG 8.OH H O E2

2Ž .

Ž .From Eq. 2 , the standard state free energy offormation of ettringite, DG 8, can be estimated usingE

the standard state free energies of its ions and thesolubility product constant ett.K . However, thesesp

equations are valid, and the thermodynamic data canbe derived from solubility measurements only whenettringite is in its standard state, dissolves congru-ently according to the reaction described above, andequilibrium is achieved during dissolution.

1.2. PreÕious studies on ettringite solubility and( ) ( )geochemistry of the Ca OH –Al SO –H O sys-2 2 4 3 2

tem

The earliest ettringite solubility study was con-Ž .ducted by Jones 1944 , who reported that ettringite

Ž .Ca-sulfoaluminate dissolves incongruently to gyp-sum and Al-hydroxide at low pH, and to portlanditeat high pH. Although Jones did not estimate ettrin-gite solubility, the aqueous ion concentrations he

( )S.C.B. Myneni et al.rChemical Geology 148 1998 1–19 3

reported were later used by Hampson and BaileyŽ .1982 to estimate the solubility product of ettringite.Using gypsum and hydrogarnet standard free ener-gies of formation, and the data reported by JonesŽ . Ž .1944 , Hampson and Bailey 1982 predicted thatett. y36 y45 ŽK would vary between 10 and 10 withspŽ .y Ž ..Al OH in Eq. 1 . They attributed this variation4

in ett.K to variations in ettringite crystal length andspŽ .habit. More recently, Atkins et al. 1991 reported a

ett. y111.3 Ž .K value of 10 for Eq. 1 . Based on thesp

molar ratios of CarSO and CarAl in ettringite and4

its saturated solution, these authors suggested thatettringite dissolves congruently above a pH of 10.7.

Ž .Damidot et al. 1992 and Warren and ReardonŽ . Žett.1994 have also obtained similar results K :sp

y44.55 y44.91 Ž .y10 and 10 , respectively, with Al OH in4Ž .. ett.Eq. 1 and showed that K is invariant withsp

Ž .changes in solution composition and pH 11.0–13.0 .However, the X-ray diffraction profiles of synthetic

Ž .ettringite prepared by Warren and Reardon 1994Žshow minor intense peaks around 11–12 82u Cu–

.K a , which do not correspond to ettringite. Suchimpurities were also noticed in some of our ettringitesynthesis when concentrated Al-sulfate solutionswere mixed with CaO in the absence of sucrose. Asdiscussed later, we used sucrose in ettringite synthe-sis and such impurities were not noticed in the finalproduct. Another value for ett.K is used by thesp

Ž .EQ3r6 computer code Sarkar et al., 1982 , whichdiffers from the reported value of Atkins et al.Ž .1991 by 3 log units. Thus, ettringite solubility isnot clear with several published values present in theliterature. The discrepancy in the reported values

Ž .may be attributed to 1 insufficient attainment ofŽ .equilibrium, 2 variations in synthetic ettringite stoi-

Ž .chiometry, and 3 probable carbonate contaminationin the synthetic ettringite examined in previous stud-ies.

( ) ( )1.3. Geochemistry of the Ca OH –Al SO –H O2 2 4 3 2

system at 298 K

Numerous investigators have examined theŽ . Ž .Ca OH –Al SO –H O system at pH ) 11.02 2 4 3 2

ŽJones, 1944; Atkins et al., 1991; Havlica and Sahu,1992; Damidot et al., 1992; Damidot and Glasser,

.1993 , but little information is available for the pHrange, 4.0 to 9.0. Since the geochemistry of CaO–

Ž .SO –H O and Al SO –H O systems is well un-3 2 2 4 3 2

derstood, available literature on these two systemswas grouped to interpret mineral interactions in the

Ž . Ž .Ca OH –Al SO –H O system in the neutral pH2 2 4 3 2

range.Ž .Jones 1944 studied the stable and metastable

Ž . Ž .mineral phases of the Ca OH –Al SO –H O2 2 4 3 2

system and effects of alkali on the stability of thesemineral phases in the alkaline pH range. He con-cluded that minerals forming in this system are solidsolutions of ettringite and monosulfoaluminateŽ Ž .Ž . .Ca Al SO OH P6H O and that these coexist4 2 4 12 2

Ž Ž . .with portlandite Ca OH , gypsum, and gibbsite2Ž .Fig. 1 . However, ettringite was found to be themost stable phase in the presence of highly alkaline

Žand sulfate-rich solutions Jones, 1944; Damidot and.Glasser, 1993 . In addition, temperature, dissolved

CO , and H O activities can strongly influence et-2 2

tringite stability. High temperatures stabilize mono-Žsulfoaluminate over ettringite Damidot and Glasser,

.1992 and high CO and low H O activity decom-2 2Ž .posed ettringite to aragonite CaCO with vaterite3

Ž . ŽCaCO as an intermediate phase Nishikawa et al.,3.1992 .

In the pH range 5.0 to 9.0, relatively few mineralsŽ . Žform in the Ca OH –SO –H O system Doner and2 3 2

.Lynn, 1989 . The known mineralogy consists ofŽ .gypsum CaSO P2H O and its analogues with dif-4 2

ferent amounts of hydration water such as bassanite

ŽFig. 1. Ettringite stability in alkaline environments Hampson and.Bailey, 1982 .


Ž .CaSO P0.5H O , and anhydrous phases including4 2Ž .anhydrite CaSO and its polymorphic forms a- and4

g-CaSO . Of these phases, gypsum is commonly4Žencountered in soils and weathered fly ashes Lu,.1981; Doner and Lynn, 1989; Fowler et al., 1993 .

Ž .The Al SO –H O system is more complicated2 4 3 2

and includes several Al-hydroxide phases such asŽ Ž . . Ž .gibbsite g-Al OH , boehmite g-AlOOH , and di-3Ž . Žaspore a-AlOOH found at near neutral pH Hsu,

.1989; Hemingway and Sposito, 1996 , and basicŽ Ž . Ž .Al-sulfates, such as basaluminite Al OH SO P4 10 4

. Ž Ž . Ž . .5H O , aluminite Al OH SO P7H O , jurban-2 2 4 4 2Ž Ž . . Ž Ž .ite Al OH SO P5H O and alunogen Al SO P4 2 2 4 3

. Ž . Ž17H O at acidic pHs -7.0 Singh and Brydon,2

1970; Adams and Rawajfih, 1977; Nordstrom, 1982;Khanna et al., 1987; Reardon, 1988; Nordstrom and

.Alpers, 1997 . Several other Al-sulfates with Na, andK in the crystal structures are also reportedŽ .Nordstrom, 1982 . Experimental studies on thechemistry, mineralogy and crystal structures of somebasic Al-sulfates were conducted by Bassett and

Ž . Ž .Goodwin 1949 , Henry and King 1950 , JohanssonŽ . Ž .1960, 1962 and Wang and Hsu 1994 , and theavailable thermodynamic data on several of these

Ž .phases was compiled by Nordstrom 1982 .In spite of extensive laboratory and field studies,

the identification of phases controlling Al solubilityin natural systems is still a subject of controversyŽ .Nordstrom and Alpers, 1997 . Soil solutions at nearneutral pH are commonly supersaturated with respectto gibbsite, and it has been proposed that microcrys-talline andror amorphous Al-hydroxide precipitationmay control Al solubility in such environmentsŽNordstrom and Ball, 1986; Sullivan et al., 1988;

.Nordstrom and Alpers, 1997 . Thermodynamic spe-ciation of acidic surface and ground waters, andmine spoil leachate are reported to be saturated with

Žbasaluminite around pH 5.0 Theobold et al., 1963;. ŽSjostrom, 1993 and jurbanite F5.0 van Breemen,¨ ¨

1976; Sullivan et al., 1988; Eary et al., 1991; Prenzel.and Shulte-Bisping, 1995 . However, neither of these

phases was identified in soils or weathered rock byŽ .direct physical methods Longmire et al., 1990 .

Although direct precipitation of basaluminite oc-curred when bentonite was reacted with Al-sulfate

Ž .solutions Singh and Brydon, 1969, 1970 , the initialprecipitates formed upon reaction of aqueous

Ž . Ž .Ca OH and Al SO were poorly crystalline2 2 4 3

basaluminite. Sample aging and heating were re-Žquired to produce crystalline basaluminite Adams

.and Hajek, 1978 . These studies clearly indicate thatthe identity of the solids controlling the activity ofaqueous Al3q and SO2y at pH values near 4.5 and4

at ambient temperatures is unresolved.The present study examines ettringite solubility,

across the pH range of 3.8 to 12.5. Direct physicalŽmethods X-ray diffraction, electron microscopy, and

.Fourier Transform infrared spectroscopy were usedto identify reaction products formed during theweathering process. This information was combinedwith mass balance data, equilibrium, and reaction-path geochemical modelling, and existing literature

Ž .data to elucidate the geochemistry of the Ca OH –2Ž .Al SO –H O system. As will be discussed, this2 4 3 2

study supports the existence of Al-hydroxy sulfatephases below neutral pH and shows that they canprecipitate rapidly in natural systems, and thus poten-tially influence major and trace element dynamics inthese environments.

2. Experimental materials and methods

2.1. Materials

2.1.1. Solutions, glassware and reagentsDeionized water produced by a Barnstead

NANOpure II system was used in all experiments.Before beginning of each experiment, this water wasboiled for an hour to drive off dissolved CO and2

cooled in a glove box filled with N gas. The2

dissolved SO2y, Cly, Ca2q, and Naq ion concentra-4

tion in these waters were below detection whentested by inductively coupled plasma emission spec-trometry and ion chromatography. Teflonw andpolypropylene bottles and centrifuge tubes were used

Ž .for experiments at high pH pH)8.0 and glasswareat low pH. All labware was washed according to the

Ž .following protocol: 1 soak 24 h in 2% MicroŽLiquid Laboratory Cleaner Cole–Parmer Instru-

. Ž . Ž .ment , 2 soak 24 h in 10% HCl, 3 wash and rinseŽ .with deionized water. Reagent grade CaO Aldrich ,

Ž . Ž . Ž .Al SO P 16H O Baker , sucrose Jenneile ,2 4 3 2Ž . Ž .NaOH Jenneile , and HCl EM Science were used

in the ettringite synthesis and dissolution experi-


ments. These chemicals were used as supplied. X-rayŽ .diffraction XRD and Fourier Transform infraredŽ .spectroscopy FTIR measurements of CaO indicated

that it was free from CaCO impurities. Reference3

solutions supplied by GFS Chemicals were used toŽ .calibrate the ion chromatograph IC , and atomic

Ž .absorption AAS and inductively coupled plasmaŽ .ICP spectrometers.

2.1.2. Ettringite synthesis and characterizationNatural ettringite is difficult to obtain in large

quantities and is usually contaminated with foreignelements such as Si, Fe, K, Na and CO . Thus, a3

Ž .synthesis described by McMurdie et al. 1986 wasadopted to obtain crystalline ettringite. In this syn-

Ž .thesis, ettringite was made by dissolving Al SO P2 4 3

16H O and CaO in the appropriate molar propor-2

tions in 10% sucrose solution. Nucleation was rapidand a suspension appeared immediately after mixing.This suspension was continuously stirred for 48 hand then the solids were separated by centrifugation.

Ž .These solids were dried in a desiccator under N2

after repeated washing with aqueous NaOH at pH12.0 to remove adsorbed sucrose and any gypsumprecipitated during synthesis. This synthesis wasconducted at room temperature and inside a glovebox filled with N gas. The precipitated solids were2

examined by XRD, FTIR, scanning electron mi-Ž . Ž .croscopy SEM and thermogravimetry TG . The

chemical composition of the precipitate was exam-Ž .ined by dissolving 0.1 g of oven-dried 383 K

material in 100 ml of 0.1 M HCl and analyzing forAl, Ca, and SO . The obtained data were normalized4

to 6 moles of Ca per mole of ettringite. Loss ofŽ . Ž .moisture wt.% at 383 K from TG analysis was

converted to moles of H O based on the reported2

stoichiometric composition of ettringite. The surfacearea of the synthetic ettringite was measured usingsingle point N -BET surface area measurements.2

2.2. Methods

2.2.1. Solubility studies, ambient pHThe ettringite solubility product was evaluated

using ettringite synthesized as described above, andequilibrium was approached from both undersatu-

Žrated and supersaturated solutions by means of the.common-ion effect . For equilibrium from undersatu-

ration, 0.25 g of ettringite was reacted with 25 mlwater in polypropylene tubes inside a glove boxfilled with N gas. To minimize electrolyte-ion sub-2

stitution in ettringite no background electrolyte wasŽused to maintain the ionic strength Kumarathasan et

.al., 1990; Pollmann et al., 1989 . The headspace in¨tubes was filled with N gas, lids were tightly2

screwed, and the tubes were later agitated continu-Ž .ously at 298 "1 K. After 0.02, 0.08, 0.17, 0.25,

0.5, 1, 4, 24, 120 and 450 h, three replicate tubeswere removed and centrifuged for 10 min at 7000rpm. The supernatant pH was measured and thesolutions were filtered through a 0.45 mm pore sizeNucleopore polycarbonate membrane filter. Thesolids were air-dried, and studied using XRD and

ŽFTIR to characterize the type of dissolution con-.gruent or incongruent . The chemical composition of

the filtered solutions and solids was determined asdescribed previously. The pH and the suspensionconcentration effects on ett.K were also evaluatedsp

in a similar way. The solution pH of these sampleswas maintained at the required values by NaOH orHCl addition. To approach equilibrium from super-saturation, ettringite saturated solutions were firstprepared by equilibrating synthetic ettringite indeionized water for a week. The solutions wereseparated from the solid phase and spiked with CaO

Ž .and Al SO P16H O solutions. Ettringite precipi-2 4 3 2

tation from these supersaturated solutions was in-duced by seeding with synthetic ettringite. The re-sulting solids were separated by filtration after 48 h,and the filtered solutions and solids were analyzed.

2.2.2. Ettringite solubility around neutral pHŽ .Laboratory synthesized ettringite 0.5 g was re-

acted with deionized water in polypropylene cen-trifuge tubes and the sample pH was adjusted by

Žadding concentrated HCl differing strength depend-.ing on the final required pH in a glove box filled

with N gas. Solution pH was adjusted between 42

and 9.5 in 0.5 pH units and later the reacted solidsŽ .were continuously agitated at 298 "1 K. Three

vials for each pH were collected after 10, 120, and240 days and analyzed as described previously.

2.2.3. Geochemical speciation and modellingSpeciation calculations were conducted for total

dissolved Ca2q, Al3q and SO2y using the4


ŽMINTEQA2 and EQ3 thermodynamic codes Al-.lison et al., 1990; Wolery and Develer, 1992 . Prior

to speciation, the thermodynamic data bases of thesecodes were modified with recently reported solubil-ity constants for various Ca and Al-sulfates, and

Ž .Al-oxy–hydroxides Appendix A . Davies equationwas used to calculate single ion activity coefficients.

ŽPotential mineral phase saturation indices SI, Vsw x.log IAPrK were estimated using the computedsp

ion activities and the reported literature values ofmineral solubility products. Reaction path calcula-tions for incongruent dissolution of ettringite wereperformed using EQ6. This program assumes thatchemical reactions are irreversible and they can berepresented by a sequence of partial equilibriumstates. Each equilibrium step is reversible with re-spect to the next, but overall the system is irre-

Ž .versible to the initial state Helgeson, 1968 . Foreach infinitesimally small step in equilibrium state,the law of mass action, mass balance, charge balanceand non-ideality are solved as the reaction progressestowards the final state. As the reaction proceeds fromthe initial stage, changes in solution composition arerepresented by the reaction progress variable, j ,which increases as equilibrium is approached. Thereaction can proceed along the considered path onlyif the free energy change of the reaction with respect

Ž .to j is negative dGrdj-0 . Since the reactionrates are not available for several reactions, theconsumed reactant molar concentration was used toestablish relative rates in this model. The presentdata were simulated as a closed system, such that theprecipitated solid phases during the reaction re-mained in the system and were allowed to change asthe reaction progressed along the reaction path. Al-though real geochemical systems are open to theinputs of foreign ions that are not considered in thepresent system, mineralogically closed models offer

Ž .heuristic value to studies of the Ca OH –2Ž .Al SO –H O system.2 4 3 2

2.3. Instrumentation and analysis

XRD, SEM and FTIR were used to examine thesolid phase mineralogy and composition. The XRDpatterns were collected using Cu–K a radiation, anda Philips PW 1216r90 wide range goniometer fittedwith a theta-compensating slit and a graphite mono-

chromator. The instrument was calibrated usingŽ .cholesterol and Si powder NIST SRM 640b for low

and high 2u angles, respectively. The samples usedfor XRD scan were finely powdered, mounted on asilicon holder, and the powder surface was smoothedwith a glass slide prior to analysis. The diffractionscans ranged from 6 to 55 82u with a 0.05 82u stepinterval and a 4 s per step counting time. The sampled-spacings were compared with reported data fromthe ICDD powder files. Diffuse reflectance FTIRspectra were collected from 4000 to 400 cmy1 with2 cmy1 resolution on a Mattson Polaris FTIR. AJEOL JSM-820 SEM was used to study the mineral-ogy at trace concentrations. Energy dispersive X-ray

Ž .analysis EDX of the solid phase elemental distribu-tion was made on the same instrument and processed

Žwith software supplied by Oxford Inst. Link Analyt-.ical eXL . Single point surface area measurements

were made with a Micromeritics Flow Sorb II 2300BET surface area analyzer using N adsorption. The2

instrument was calibrated with specimen kaolinŽ 2 y1.standard a8570, surface areas10.3 m g and

Ž 2alumina standard a8571, surface areas153 my1 .g supplied by National Bureau of standards.

Aqueous phase cationic concentrations were ana-lyzed with a Perkin-Elmer 3030B AAS. A Dionex

Ž .2000i Basic Chromatography Module IC , a AS4A4 mm separator column, and an AG4A guard columnwere used to measure anion concentration. Thechemical analyses obtained with the above tech-niques were confirmed with a Leeman PS2000 ICPspectrometer.

2.3.1. pH measurementAn EA 920 Orion Expandable Ion Analyzer con-

Ž .nected to a Ross pH electrode 8103 was used tomeasure solution pH. This instrument was calibratedwith pH 10.0, 7.0, and 4.0 buffers supplied by VWRScientific. The buffer solutions were stored in an

Ž .incubator at 298 "1 K and transferred to a beakerbefore each electrode calibration. To minimize liq-uid-junction potential effects, samples were first cen-trifuged for 15 min at 7000 rpm and pH measure-ments were made later in particle-free supernatants.The pH readings usually stabilized in 4 min, and themeasurements were taken after that period. The sam-ple and reference solution pH measurements werealso repeated routinely to check for any drift in pH.


3. Results and discussion

3.1. Characterization of synthetic precipitate

The solids prepared in the synthesis were exam-ined by XRD, FTIR, SEM and TG. The solid XRDpatterns were equivalent to the data reported in the

Ž .International Center for Diffraction Data ICDDmanual. This manual uses different references for

Žsynthetic ettringite XRD carda: 31-251, 9-414, 37-.1476 and 41-1451 . While the line d-spacings in

these references match, their intensities are not iden-tical. Similarly, the synthetic ettringite d-spacingsalso matched all previously reported data but theintensities were similar to those reported in carda

Ž . Ž .37-1476 McMurdie et al., 1986 Fig. 2 .The IR spectra of the synthetic material compare

Ž .well with those reported by Pollmann et al. 1989¨Ž .and Kumarathasan et al. 1990 . The vibrational

modes of SO2y and H O molecules in synthetic4 2

ettringite agree with ettringite crystal structure re-Ž .finements Moore and Taylor, 1970; Myneni, 1995 .

A discussion of these is not appropriate here, and theŽ .interested readers may refer to Myneni 1995 . The

absence of CO2y vibrational frequencies around ;3

2500–3000, 1400, and 750 cmy1 indicated that the

synthetic material was initially free from calciteŽ .White, 1974 . However, upon long exposure to the

Ž .atmosphere ;60 days , a doublet appeared aroundŽ . y11450 1488, 1430 cm due to CO adsorption. On2

the basis of these XRD and FTIR results, the syn-thetic material was identified as crystalline ettringite.The synthetic ettringite crystals were needle shaped

Ž .and 2 mm long from SEM , and had a surface area2 y1 Ž .of 7.9 m g from BET surface area analysis

Ž .Myneni, 1995 .The water content of the synthetic ettringite as

calculated from weight loss at 383 K, correspondedto a stoichiometry of 26 moles of H O per formula2

unit. Since this number includes both the adsorbedand structural H O, the absolute quantity of struc-2

tural H O remains uncertain. Fortunately, this num-2

ber does not affect the solubility product as long asH O activity is near unity. Wet chemical analysis of2

Ž .synthetic ettringite normalized to 6 moles of Caindicate that its chem ical form ula was

Ž . Ž . Ž .Ca Al SO OH P26H O Table 1 . Al-6 2.02 4 2.79 12.48 2

though the reported stoichiometry based on this wetchemical analysis and the unit cell were similar,minor variations in Al, sulfate and OH molar concen-tration exist. Such differences may also have beendue to OH substitution for SO2y in the tunnels of4

Fig. 2. XRD of synthetic ettringite.


Table 1Chemical analysis of synthetic ettringites

Sample m m m mCa Al SO OH4

ETT1 6 2.03 2.76 12.57ETT2 6 2.02 2.80 12.46ETT3 6 2.02 2.79 12.48ETT4 6 2.03 2.80 12.49ETT5 6 2.05 2.73 12.69ETT6 6 2.02 2.79 12.48

Samples marked ETT1-2 are synthesized as reported by McMur-Ž .die et al. 1986 . Samples ETT3-4 and ETT5-6 are also synthe-

Ž .sized in a similar way but with excess CaO, and Al SO ,2 4 3

respectively. Data is reported as number of moles of an elementper one mole of ettringite and are normalized to 6 moles of Ca.

Ž .ettringite Pollmann et al., 1989 . Similar molar pro-¨portions and non-stoichiometry have been previously

Ž .reported for ettringite Hassett et al., 1991 . Thecomposition of ettringite was also found to be invari-ant with changes in the initial ratios of the reactantsŽ Ž . Ž . . Ž .Ca OH and Al SO Table 1 . However, the2 2 4 3

Ž .precipitate crystal size was longer 6–12 mm whenthe syntheses were conducted in the absence of

Ž .sucrose and at higher solution pH Myneni, 1995 .

3.2. Ettringite solubility product

Ettringite dissolution in undersaturated solutionswas rapid with reaction reaching equilibrium within

Ž .minutes Table 2; Fig. 3 . In contrast, approach tothe equilibrium from supersaturated solutions wasrelatively slow and took about 24 h. However, awhite ettringite precipitate appeared immediately af-

Ž .ter CaO and Al SO solutions were mixed, and2 4 3

the IAP was close to the ett.K after a few minutessp

of mixing. Slow precipitation kinetics have beenobserved for several mineral systems when the IAP

Žis close to the K Stumm, 1993; Sposito, 1994;sp.Oelkers et al., 1994 .

The CarAl and CarSO ratios for synthetic et-4Žtringite were 2.97 and 2.22, respectively Tables 1

.and 3 . Little or no change was observed for these

Table 2Žett. .Solubility product constant K measurement for ettringitesp

aUndersaturated solutions

Ž . Ž .Time pH Ionic strength Total concentrations p ion pIAP Avg. pIAP p IAP 33 y3Ž . Ž . Ž . Ž .h "0.1 mol my mol m

3q 2q 2y 3q 2q 2yAl Ca SO Al Ca SO4 4

0.02 11.0 6.72 0.78 2.21 1.26 24.45 2.86 3.13 111.4 111.8 112.90.08 11.1 7.95 0.87 2.35 1.24 24.79 2.84 3.15 110.9 110.8 112.30.17 11.0 7.35 0.90 2.12 1.12 24.38 2.87 3.18 111.5 111.6 113.00.25 11.2 7.35 0.77 2.13 1.11 25.25 2.87 3.18 110.9 111.0 112.30.5 11.2 8.27 0.77 2.11 1.08 25.25 2.88 3.19 110.9 111.0 112.31 11.3 8.58 0.66 2.01 0.97 25.72 2.70 3.24 109.7 111.0 111.14 11.1 7.48 0.73 2.05 1.02 24.87 2.88 3.22 111.5 111.4 112.924 11.0 7.43 0.72 2.00 1.08 24.48 2.89 3.19 111.9 112.1 113.4120 11.0 8.41 0.75 2.12 1.05 24.68 2.71 3.05 110.8 111.9 112.3450 11.0 8.06 0.79 2.08 0.92 24.68 2.72 3.10 111.0 111.0 112.5

Supersaturated solutions24 12.4 8.97 0.02 14.7 0.28 31.71 2.26 3.73 107.4 108.0 108.124 12.2 9.46 0.11 13.7 0.23 30.16 2.24 4.13 107.7 107.7 108.548 11.5 8.47 0.17 3.73 0.29 27.21 2.63 3.83 111.5 112.0 112.648 11.4 7.98 0.17 3.03 0.45 26.72 2.71 3.61 111.7 111.7 113.0

a Equilibrium from undersaturated solutions: synthetic ettringite was dissolved in deionized water, and samples were collected for differentŽ .reaction periods. These samples were analyzed in triplicate and the average ylog IAP is given as avg. pIAP. For each reaction time,

Ž Ž ..complete chemical analysis of one of the triplicates and its pIAP according to Eq. 1 are given. For the same samples, IAP was alsoŽ Ž .. Ž . y1calculated using the measured stoichiometry according to Eq. 3 and reported as p IAP . Suspension concentration was 10 g l .3

Equilibrium from supersaturated solutions: ettringite saturated solutions were prepared initially by dissolving synthetic ettringite inŽ . Ž .deionized water, and ettringite was allowed to precipitate from these solutions by adding excess dissolved Ca OH and Al SO .2 2 4 3

Ž . Ž .The ion activities were computed using the MINTEQA2 computer code and reported as p ion . Ts298 "1 K.


Fig. 3. Ettringite solubility in deionized water.

Želemental ratios of the solids after dissolution Table.3 . In addition, similar CarAl and CarSO molar4

ratios were observed for both the solid and solutionphases. Taken as a whole, these data support congru-ent dissolution of ettringite as proposed by Atkins et

Ž .al. 1991 . Further support for congruent dissolutionwas also provided by XRD and FTIR which showedno effect of dissolution on the spectral properties ofettringite for all samples with solution pH values

Žabove 10.7. Thus, in this hyperalkaline regime pH.s12.5 to 10.7 , ettringite dissolution can be de-Ž .scribed by Eq. 1 . At pH values below 10.7, incon-

gruent ettringite dissolution produced gypsum andamorphous Al-hydroxides. This will be discussed inSection 3.3.

3.2.1. Effect of stoichiometry on et t.K s p

Variations in crystal stoichiometry can result indifferent ett.K values. The stoichiometry estimationsp

based on the wet chemical analysis uses normaliza-tion procedures and thus deviates from the unit-cellderived crystal formula. In the case of ettringite, themeasured stoichiometry was close to the ideal stoi-

Ž Ž ..chiometry as described in Eq. 1 , but with slight2y Ždeviation for OH and SO molar composition Ta-4

.ble 1 . Based on this measured stoichiometry, ettrin-gite dissolution can be expressed as:

Ca Al SO OH P26H Om6Ca2qŽ . Ž . 12 .486 2 .02 4 22 .79

q2.02Al3qq2.79SO2y q12.48OHyq26H O4 2

3Ž .and the solubility product constant for this reactionŽ . y112.9 ett.K is 10 . Thus, K and K are ettringite3 sp 3

solubility product constants estimated from the idealand measured stoichiometries, respectively. Al-though, the average K is about one log unit smaller3

than the average ett.K , these two values are withinsp

the standard error of the chemical analyses.

3.2.2. Effect of pH on et t.K s p

The influence of pH on the solubility productconstant was evaluated over the ettringite pH stabil-

Ž .ity range 10.7–12.5 . The solubility studies con-Žducted at pH 10.7, 11.3, 12.2, and 12.5 at constant

y1 .suspension concentration, 4 g l produced thesame ett.K values, suggesting little pH influence onsp

Table 3Ž . Ž .Chemical analysis of synthetic ettringite before ETT 2 , and after solubility measurement rest of the samples

Ž .Time h Solid Equilibrium solution

Moles of each component Molar ratios Molar ratios

Ca Al SO OH CarAl CarSO CarAl CarSO4 4 4

ETT 2 6 2.02 2.80 12.47 2.97 2.14 y y0.17 6 2.00 2.72 12.57 3.0 2.21 2.35 1.904 6 2.04 2.67 12.80 2.94 2.25 2.79 2.0124 6 2.01 2.70 12.64 2.98 2.22 2.80 1.85120 6 2.01 2.67 12.69 2.98 2.25 2.82 2.02450 6 2.03 2.81 12.47 2.96 2.14 2.62 2.25720 6 2.02 2.70 12.67 2.97 2.22 2.84 2.04

All elements of the solid are reported as the number of moles per one mole of ettringite and are normalized to 6 moles of Ca.


Table 4Effect of pH and suspension concentration on ettringite solubility

aa. Effect of pHet t.Ž .pH p ion p Ksp

2q 2y 3qŽ ."0.1 Ca SO Al4

10.7 2.62 2.59 22.5 109.311.3 2.79 3.12 26.31 111.112.2 3.21 3.32 29.24 109.812.3 2.31 4.02 31.81 110.0

bb. Effect of suspension concentrationett.Ž .Suspension pH p ion p Ksp

y1Ž . Ž .density g l "0.1 2q 2y 3qCa SO Al4

4 11.4 2.90 3.14 25.80 110.110 11.1 2.84 3.06 24.49 110.312.1 11.0 2.78 3.00 25.30 112.420 10.8 2.85 2.92 23.44 111.030 10.5 2.66 2.84 22.14 110.8

Žett. .The solubility product constant K was calculated based onspŽ .Reaction 1. Ion activities are reported as p ion .

a y1 Ž .Suspension concentration: 4 g l . T s298 "1 K.b Ž .Equilibrium pH was allowed to change. T s298 "1 K.

ett. Ž .the K Table 4a . However, high pH samplessp

exhibit high dissolved Al concentrations with nosignificant trend for Ca2q and SO2y activity. Simi-4

lar observations were also reported by Atkins et al.Ž . Ž .1991 , Damidot et al. 1992 , and Warren and Rear-

Ž .don 1994 .

3.2.3. Effect of suspension concentration on et t.K s p

Laboratory mineral solubility-product-constantstudies frequently employ unusually low solid-to-solution ratios which suppress secondary mineralformation. Natural systems, on the contrary, mayexhibit extremely high solid-to-solution ratios. Whilethe suspension concentration should not affect ett.K ,sp

it may have a strong effect on its dissolution rate andthus may significantly control weathering reactionsŽ .Casey et al., 1993 . Suspension concentration alsomodifies dissolution patterns for minerals with sev-

Ž 2yeral chemical moieties in their unit-cell e.g., SO ,4y .OH , etc. . Different functional groups may exhibit

contrasting release-rates and this facilitates sec-ondary, meta-stable phase formation. To examinethis behavior, ettringite dissolution experiments wereconducted at suspension concentrations of 4, 10,12.1, 20, and 30 g ly1. The results indicate that

ett.K was invariant at all the suspension concentra-sp

tions examined and equilibrium was achieved withinminutes for all cases. However, as the solid-to-solu-tion ratio increased, pH dropped and the aqueous

2q 3q 2y ŽCa , Al , and SO concentration increased Ta-4.ble 4b . In addition, at high suspension concentration

Ž y1 .)10 g l the equilibrium solution CarAl, andŽ .CarSO molar ratios 2.3 and 1.7, respectively4

were lower than those of the reacting synthetic et-Ž .tringite 2.97 and 2.22, respectively . This discrep-

ancy may be attributed to the formation of surfacepolymers or a secondary precipitate at high suspen-

Žsion concentration Wollast and Chou, 1992; Casey.et al., 1993 . Thus, these secondary phases and

ettringite are at equilibrium with the solution phaseat high solid concentrations. Further studies are re-quired to examine the nature of these precipitates.

3.2.4. Standard state free energy of formation ofettringite

Based on the measured ett.K and the free ener-sp

gies of formation of Ca2q, Al3q, SO2y, and H O4 2Ž .Appendix A , the standard state free energy offormation for ettringite was estimated as y15 204.7

y1 Ž y1 ."23 kJ mol y3634.0 kcal mol . This value isŽ .closer to the value reported by Atkins et al. 1991

Ž y1y9048.2 kJ mol , without 26 molecules of H O2.in ettringite , and the estimated value based on the

solubility product constant measured by Warren andŽ . Ž y1 .Reardon 1994 y15 207.0 kJ mol .

3.3. Ettringite incongruent dissolution and geochem-( ) ( )istry of the Ca OH –Al SO –H O system2 2 4 3 2

3.3.1. Solution chemistryThe chemical composition of aqueous solutions

reacted with ettringite were dependent on equilib-rium pH, but no reaction time effect was evident for

Ž .the period of 10 to 240 days Table 5; Fig. 4 . Thismay suggest that the equilibrium was achieved inF10 days. Around pH values of 9.0 and 4.5, the

Žaqueous ion concentration changed dramatically Ta-.ble 5 . As the pH was lowered from 10, a threshold

was crossed at pH;9.0. At this point, dissolvedCa2q and SO2y concentration increased from 134

mM to 90 mM, and 5.6 mM to 10 mM, respectively,while the concentration of dissolved Al3q decreasedfrom 0.2 mM to 0.03 mM. As the pH was lowered


Table 5Ettringite incongruent dissolution at pHF10.0

Ž . Ž .pH "0.1 Equilibrium Total aqueous ion concentrations mM Identified equilibrium solidsŽ . Ž .time days xsXRD, ssChemical speciation2q 3q 2yCa Al SO4

3.8 10 71.28 23.07 16.03 jurbanite , basaluminite , gypsums s x,s

120 69.58 23.15 15.78240 70.75 21.33 14.93

4.2 10 86.8 2.63 10.09 jurbanite , basaluminite , gypsums s s, x

120 85.65 3.22 10.09240 85.45 3.22 10

4.9 10 93.05 0.06 9.53 basaluminite , Al-hydroxide , Gypsums s s, x

120 89.43 0.14 9.47240 92.75 0.07 9.31


240 90.28 0.03 9.476.0 10 90 0.03 9.77 basaluminite , Al-hydroxide , Gypsums s s, x

240 89.65 0.1 9.136.8 240 93.93 0.02 9.49 basaluminite , Al-hydroxide , Gypsums s s, x


8.0 10 93.95 0.02 10.09 Al-hydroxide , Gypsums s, x

120 87.4 0.02 9.63240 89.83 0.11 9.66

8.5 120 73.98 0.06 9.63 Al-hydroxide , Gypsums s, x

240 77.33 0.06 9.949.4 10 41.25 0.03 10.06 Al-hydroxide , Gypsum , Ettringites s, x x

120 43.05 0.22 10.31240 43.7 0.24 10.34

10.0 10 13.65 0.22 5.84 Al-hydroxide , Gypsum , Ettringites s, x x

240 13.4 0.34 5.59

Ž .from the first threshold to the second threshold 4.5 ,the concentrations of dissolved Ca2q, Al3q and SO2y

4

were invariant. Below a pH of 4.5, Ca2q concentra-tion decreased from 90 mM to 70 mM, while that of

3q 2y Ž .Al and SO increased sharply Table 5 . The4

activities of Ca2q and SO2y showed the same trend4Ž .as their total concentration Fig. 4a and b . The

decrease in Ca2q activity below pH 5.0 is notewor-thy and will be addressed below. In the experimentalpH range, Al3q activity exhibited several inflectionpoints which correlate with the first, second and third

Ž . ŽAl hydrolysis Fig. 4c Faure, 1989; Stumm and.Morgan, 1994 .

( )3.3.2. Saturation index SI calculationsDue to the propensity of amorphous and micro-

crystalline solids to form in natural weathering envi-ronments, MINTEQA2 saturation indices were calcu-lated with the precipitation of macrocrystalline solidsŽ .gibbsite, bayerite, diaspore, boehmite prohibited

ŽLindsay, 1979; Stumm, 1993; Nordstrom and Alpers,.1997 . In the case of basaluminite, the lack of solu-

bility estimates for poorly-crystalline phases necessi-tated the use of crystalline solids in saturation in-dices calculations. The SI calculations showed thatonly ettringite was stable above pH 10.7 and thatgypsum and amorphous Al-hydroxides precipitate

Ž .between 10.5 and 4.5 Table 5; Fig. 5 . In addition,Ž . Ž .basaluminite pH-7.5 and jurbanite pH-5 were

found to be the potential Al-sulfate phases belowneutral pH. Basaluminite was always found to besupersaturated by several orders of magnitude over

Ž .the pH range of 4.0 to 8.0 Fig. 5 . Indeed, SI valuesfor this mineral increased from 0.25 at pH 3.8 to amaximum of 7.0 at pH 6.0, and then decreased to 0.3at pH 7.7.

3.3.3. Reaction path calculationsNon-equilibrium reaction path calculations were

conducted to simulate acid titration of ettringite from


Fig. 4. Ettringite dissolution at pHF10. The ordinate indicatesion activities.

Fig. 5. Ettringite dissolution at pHF10. Saturation indices forŽ . Ž .different minerals in Ca OH –Al SO –H O ternary system2 2 4 3 2

were calculated using the solubility product constants derivedfrom the Gibbs free energies given in Appendix A.

its ambient equilibrium pH to a value of 4.0, and theŽprecipitation of crystalline Al-hydroxide phases e.g.,

.gibbsite inhibited. The calculations indicated thatettringite dissolved incongruently to gypsum andAl-hydroxide below pH 10.7. Basaluminite appearedas a stable phase when the pH reached 7.5 andpersisted even at low pH. Amorphous Al-hydroxidedisappeared at pH 4.5 and jurbanite formed at its

Ž .expense as pH further decreased Fig. 6 . These werein agreement with the actual mineralogical and geo-chemical observations obtained from the laboratoryweathering experiments described above. However,when the pH was -5 the predicted ion activities

Ž .were smaller 1 log unit than the experimentalvalues. Thus the stability of gypsum, jurbanite andbasaluminite alone may not adequately explain thisCa2q, Al3q and SO2y ion behavior. This suggests4

Ž .that some other phase s not accounted for in the

Ž . Ž .Fig. 6. Al-mineral equilibria in Ca OH –Al SO –H O sys-2 2 4 3 2

tem. Open circles are experimental data points and the solid linewith arrow represents the reaction path predicted using EQ6Ž 2q 2y y2.3ettringite titration with HCl; Ca and SO activites are 104

y2 .9 . 3qand 10 , respectively . Estimated Al activity in fly ashŽleachate and acidic mine spoils are also plotted B: Sullivan et al.

Ž . Ž . Ž . Ž .1988 ; l: Fruchter et al. 1990 ; % fly ashes and ' FGD :Ž ..Mattigod et al. 1990 . Minerals shown at different pH values

were identified with SEM.


thermodynamic database may also affect these ionactivities, or that the reaction path calculations didnot fully account for kinetic constraints to the forma-tion of reaction products.

3.3.4. XRD and microscopy studiesXRD analyses of the solids from the weathering

studies indicated the presence of ettringite at pH)

10.7, gypsum and ettringite between pH 9 and 10,Ž .and only gypsum below pH 9 Fig. 7 . No other

mineral phases were evident from the XRD. Thepresence of ettringite below pH 10.7 is contradictoryto the observations made from the solubility mea-surements and SI calculations. Although slower dis-solution kinetics can cause this anomaly, this is notlikely in the present system since the experimentaldata from 10 and 240 days gave similar results. Thisextended ettringite stability into the lower pH rangemay be due to its incongruent dissolution to gypsumand Al-hydroxides. The failure of thermodynamicmodels to predict the presence of ettringite below pH10.7 is presumably due to the precipitation ofmetastable or poorly crystalline phases in the systemduring incongruent dissolution.

Trace mineral identification was attempted withSEM, based on crystal morphology and energy dis-

Fig. 7. XRD of ettringite incongruent dissolution products. Peaksaround 9.3 and 11.882Q correspond to ettringite and gypsum,respectively.

persive X-ray elemental analysis. SEM revealed threemorphologically distinct euhedral and anhedral crys-

Ž .tals, and microcrystalline grains Fig. 8a . The euhe-dral crystals were monoclinic, and range in size from-1 mm to as large as 60 mm. These crystals werethe dominant phase in all the samples equilibratedbelow pH 9.0, and constitute )90% of the sample.The anhedral crystals were found in samples ofpH-7.0, ranging in size from 20 to 130 mm, and

Ž .had irregular, conchoidal-type fractures Fig. 8b .The fine-grained material was of submicron size andwas found in the pH range 8 to 5.5. The abundanceof this fine-grained fraction decreased dramaticallybelow pH 5.0.

EDX analysis indicated that the euhedral grainscontained equimolar Ca and S, suggesting that they

Ž .were gypsum crystals Fig. 9 . Some of these gyp-Žsum crystals also exhibited high Al contents 0.2 to

Ž . .0.8 "0.03 atomic percentage . Previous gypsummineralogical studies have shown little variation inchemical composition, and Al substitution has not

Žbeen reported before Deer et al., 1966; Doner and.Lynn, 1989 . While Al substitution into gypsum

cannot be ruled out, it seems likely that the measuredAl fluorescence was due to adsorbed or fine nano-scale Al-hydroxide coatings on gypsum. The an-hedral crystals are identified as Al-hydroxy sulfate

Žphases. These grains had variable AlrS ratios 1.2 to.2.6 and contained no Ca. No apparent correlation

was observed between the AlrS ratio and the grainsize of these minerals. In comparison, aluminite andbasalumanite have AlrS ratios of 2 and 4, respec-tively. Compositionally similar grains were also ob-served in prismatic form but were few in number.Morphologically similar, Al-hydroxy sulfates havebeen previously synthesized by mixing Na SO with2 4

Ž .Al-hydroxide solutions Wang and Hsu, 1994 . How-ever, no grains with composition similar to jurbanitewere detected by EDX analysis.

The fine grained material showed two distinctlyŽdifferent compositions, with Al and O in one plate-

. Ž .shaped grains Fig. 8a , and Ca, Al, S, and O in theŽ .other Fig. 8c . These two phases are likely Al-hy-

droxide and Ca–Al-hydroxy sulfate, respectively.However, it should be noted that EDX cannot distin-guish between the adsorbed and structural atoms.Thus, it is impossible to rule out the presence ofadsorbed Ca on Al-hydroxy sulfate or adsorbed Al


Fig. 8. SEM of ettringite icongruent dissolution products. a1 and a2: euhedral gypsum crystals with coatings of microcrystallineAl-hydroxides. b: Al-hydroxy sulfates. c: Ca–Al-hydroxy sulfates.

on gypsum. Nevertheless, the high Ca and Al atomicŽ .percentages 0.5–4.0 in these materials suggest that

these ions were present in structural sites rather thanas adsorbed species. To our knowledge, Ca–Al–S–Ž .OH,O phases have not been reported before in lowpH materials, such as acidic soil horizons; however,specimens of a Ca-substituted alunite, minamiiteŽ Ž . Ž . .MAl SO OH , where MsNa, K, Ca have3 4 2 6

been identified from hydrothermally altered quartz-Ž .bearing andesites Ossaka et al., 1982 .

( ) ( )3.3.5. Geochemistry of the Ca OH –Al SO –H O2 2 4 3 2

systemŽ . Ž .Several mineral phases of Ca OH –Al SO –2 2 4 3

H O system were plotted in activity–pH diagrams2

using the available mineral thermodynamic data ofŽ . Ž .the Ca OH –SO –H O and Al SO –H O sys-2 3 2 2 4 3 2

Ž .tems Fig. 5 . The minerals considered were ettrin-Ž .gite and gypsum in the Ca OH –SO –H O system,2 3 2

and amorphous Al-hydroxide, jurbanite and basalu-Ž .minite in the Al SO –H O system. Although,2 4 3 2

Ž .Wang and Hsu 1994 reported several new Al-hy-droxy sulfates, lack of their thermodynamic dataprevents us from considering them here.

The incongruent dissolution of ettringite exhibitedŽ .two pH thresholds pH;9.0 and 4.5 where the

Ž .aqueous ion activities changed dramatically Fig. 4 .Ž .At the first pH threshold pH;9.0 , gypsum, ettrin-

gite, and Al-hydroxide control the Ca2q, Al3q, and2y Ž .SO activities Fig. 4a,b,cFig. 5 . Changes in ion4

activities at these pH values correlated with the onsetof gypsum and Al-hydroxide precipitation from et-tringite. Although precipitation of these phases be-

Žgins at pH 10.7 i.e., where incongruent dissolution


ŽFig. 9. EDX elemental analysis of different mineral phases size-. Ž50 mm formed from ettringite incongruent dissolution pH range:

.4.0 to 6.0 . a: Al-hydroxy sulfate, b: Al-hydroxides, c and d:Ca–Al-hydroxy sulfates, and e: gypsum. The samples were coatedwith graphite prior to EDX measurements.

.of ettringite begins , these minerals do not signifi-cantly influence solution composition above pH 9.5.

ŽBetween the first and second pH thresholds pH.9.5–4.5 , gypsum, Al-hydroxide and Al-hydroxy sul-

fates control the Ca2q, Al3q, and SO2y activities.4

Although SO2y activity correlates to that of gyp-4

sum, calculations showed that the basic Al-sulfatesalso produce similar SO2y activities, especially be-4

Ž .low pH 7.0 Figs. 4 and 5 . In addition, SEMindicated the presence of different Al-hydroxy sul-fate phases below this pH. This suggests that thebasic Al-hydroxy sulfate phases play a critical role incontrolling the ion activities below pH 7.0. However,

Ž .the geochemistry of Ca OH –SO –H O and2 3 2Ž .Al SO –H O systems may be independent of2 4 3 2

each other in the pH range of 7 to 9.5.The onset of Al-hydroxy sulfate precipitation, at

the expense of Al-hydroxide, perturbed theŽ .Ca OH –SO –H O system and this effect is dis-2 3 2

tinct when pHF5.0. The high solubility of the basicAl-sulfates gives rise to high Al3q and SO2y activi-4

ties which, in turn, decreases aqueous Ca2q activityin equilibrium with gypsum and Al-hydroxy sulfates.This may explain the experimentally observed reduc-

tion in Ca2q activity and corresponding increase in2y Ž .SO activity below this pH Fig. 4 . In addition,4

the low Ca2q activities below pH 5.0 may also havebeen caused by the precipitation of the Ca–Al-hy-droxy sulfate phases, which were also tentativelyidentified in the SEM studies.

Equilibrium geochemical modelling withMINTEQA2 showed that basaluminite was supersat-urated by several orders of magnitude between pH 4to 8. Aluminum activity in equilibrium with basalu-minite would be smaller than was observed in theseweathering experiments. This kind of supersaturation

Ž .is caused either by: 1 slow precipitation kinetics,Ž .2 presence of basic Al-sulfates with different com-

Ž .position, andror 3 by the precipitation of an amor-Ž .phous precursor Sposito, 1994 , for example, those

tentatively identified in the SEM studies.The mineralogical relations, and the variation in

Al3q activity during ettringite incongruent dissolu-Ž 3q. Ž .tion are summarized in a log Al –pH plot Fig. 6 .

The reaction path calculations conducted for acidtitration of ettringite also predicted Al3q activity,and mineral precipitation patterns accurately abovepH 5.0. Interestingly, the Al3q activity and the

Žmineral assemblages tentatively identified basalumi-.nite, microcrystalline Al-hydroxides in the pH range

of 4.0–7.0 follow the reaction,

Al OH q3Al3qqSO2y q5H O™Ž . 3 4 2

Al OH SO .5H Oq7Hq 4Ž . Ž .104 4 2

Although the equilibrium reaction between basalumi-nite and gibbsite can be written as,

4Al OH qSO2y q2HqmŽ . 3 4

Al OH SO .5H Oq2H O 5Ž . Ž .104 4 2 2

this reaction may not occur for kinetic reasons in thetime span of few months. If this reaction controlsAl3q activity, the same Al-hydroxy sulfate phasesshould not occur at different pH and at a constantSO2y activity. This may suggest that metastable4

equilibrium exists between Al-hydroxide and Al-hy-Ž .droxy sulfate phases Reaction 4 . The measured

3q ŽAl activities of acid mine drainage Sullivan et al.,. Ž1988 , flyash Fruchter et al., 1990; Mattigod et al.,. Ž .1990 and FGD leachates Mattigod et al., 1990 also


follow the same trend, and are in support of thishypothesis. The observed similarities between theexperimental values and the field data indicate the

Ž .robustness of the predicted equilibria Fig. 6 . Theirmarked disagreement at pH-5.0 can be attributedto the formation of Ca–Al-hydroxy sulfates in thelaboratory studies and the relatively Ca poor materi-als in the field sites.

( ) ( )3.4. Geochemistry of the Ca OH –Al SO –H O2 2 4 3 23q 2 q q ( )0system: open to Fe , Mg , K , Si OH , and4

CO 2 y3

With additional ion inputs such as Fe3q, Mg2q,q Ž .0 2y Ž . Ž .K , Si OH , and CO , the Ca OH –Al SO –4 3 2 2 4 3

H O system will exhibit different mineral phases. To2

understand these effects, equilibrium speciation mod-elling and reaction path calculations were conductedat different pHs and for a range of foreign ionconcentrations. During these computations, Ca2q,Al3q, and SO2y activities were fixed at the values4

measured from the weathering experiments.Ž . Ž .When the Ca OH –Al SO –H O system is2 2 4 3 2

open to atmospheric CO , its solutions approach the2

saturation of carbonate mineral phases such as cal-cite and aragonite under alkaline pH conditions.However, the slow precipitation kinetics of calciteassociated with the slow transport of CO into the2

waste materials delays its formation relative to ettrin-gite and gypsum. Moreover, high SO2y activity4

Žfavors the formation of gypsum over calcite Doner.and Lynn, 1989 . This has also been observed from

field studies of ettringite-bearing, weathered, FGDmaterials. In these alkaline materials, calcite did not

Žappear until 240 days of weathering Fowler et al.,.1993 , whereas ettringite and gypsum precipitated

within one week of their exposure to natural weath-ering conditions. Thus, carbonate equilibria has littleinfluence on this system during the initial stages ofweathering. However, elevations in CO partial pres-2

sures from the decomposition of organic matter insubsurface soil horizons, and prolonged leaching ofSO2y from waste materials favour calcite formation.4

Speciation codes indicated that the presence of Mg2q

Ž Ž . .resulted in the formation dolomite CaMg CO ,3 2Ž . Ž Ž . .magnesite MgCO , and huntite Mg Ca CO in3 3 3 4

addition to calcite. Although huntite formation hasbeen observed in weathered dolomite or magnesite

Ž .bearing rocks Deer et al., 1966 , its rate of forma-tion has not been evaluated. Addition of Fe3q, Mg2q,

q Ž .0K , and Si OH to the system formed Fe oxyhy-4

droxides, minerals belonging to alunite–jarosite fam-ily, clays and leonhardite in addition to the other

Ž .above mentioned mineralogy of the Ca OH –2Ž .Al SO –H O system. Of the Si phases, zeolites2 4 3 2

Ž .leonhardite , and some clays such as kaolinite andhalloysite may be dominant minerals in the system.However, actual experiments are necessary to sup-port these predicted mineral equilibria.

4. Conclusions

The results presented here demonstrate that ettrin-gite dissolves congruently above pH 10.7 with a

ett. Ž . Ž .p K of 111.6 "0.8 according to Eq. 1 . Thesp

measured crystal stoichiometry, after normalizationto six moles of Ca, is slightly different from that ofthe idealized formula, but it did not affect the ett.K sp

value significantly. In addition, changes in pH andsuspension concentration had no effect on ett.K .sp

Ettringite can also exist at pH values -10.7, butonly in association with gypsum and Al-hydroxide.Ettringite completely dissolved at near neutral pHand precipitated gypsum, Al-hydroxides, Al-hydroxysulfates such as basaluminite and other previouslyunidentified Ca–Al-hydroxy sulfate phases. Al-

Ž .though the geochemistry of the Ca OH –2Ž .Al SO –H O system is simple above pH 7.0 and2 4 3 2

Ž .can be described as an association of the Ca OH –2Ž .SO –H O and Al SO –H O systems, it is com-3 2 2 4 3 2

plicated by the presence of several phases at nearneutral pH. Its simplicity above pH 7.0 arises fromthe absence of Al-hydroxy sulfate phases. EQ3r6calculations accurately predicted the ettringite weath-ering sequence and relevant ion activities when thepH was )5.0. The apparent disagreement betweenthe predicted and the observed ion activities mayhave been due to the precipitation of previouslyunknown, low pH Ca–Al-hydroxy sulfate phases,which were confirmed by the microscopic studies.Further study should explore the occurrence of thesemineral phases in field weathering environments andtheir potential to form in the presence of other ionsnot considered in the present study.


Acknowledgements

The authors would like to thank Drs. J.M. Bighamand G. Faure for their suggestions; and R.K. Fowler,U.I. Soto, and D. Beak for their assistance in theexperiments. Drs. D.K. Nordstrom, J.I. Drever, andE.H. Oelkers are greatly appreciated for providingvaluable comments and critically reviewing themanuscript. This research was supported by a grantfrom the USGS-Water Resources Program. The prin-cipal author was supported by the Ernest OrlandoLawrence Berkeley National Laboratory, Berkeley,CA during the preparation of the manuscript. Addi-tional salary and research support were provided bythe Ohio Coal Development Office, the Ohio StateUniversity, and the Ohio Agricultural Research and

Ž .Development Center OARDC . OARDC journal ar-ticle no. 49–98.

Appendix A. Standard state free energies of se-lected aqueous and mineral species

Species D G8 kJ moly1f

Aqueous species3q aAl y489.8

2q aŽ .Al OH y701.4q aŽ .Al OH y903.42

bŽ .Al OH 8 y1100.73y aŽ .Al OH y1305.04

dH O y237.22y dOH y157.3

2q dCa y553.122y dSO y744.54

Solid speciesaŽ . Ž .Al OH gibbsite y1154.893

cŽ . Ž .Al OH SO jurbanite y1490.14cŽ . Ž .Al OH SO basaluminite y4927.14 10 4dŽ .CaSO P2H O gypsum y1797.64 2

eŽ . Ž .Ca Al SO OH P y15 204.76 2 4 3 12Ž .26H O ettringite2

a Ž .References: Hemingway and Sposito 1996 ;b Ž . c Ž . d Ž .Anderson 1996 ; Nordstrom 1982 ; Faure 1989 ;eThis study.

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ettringite solubility and geochemistry of the ca oh –al so –h o...

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