Gibbsite solubility and thermodynamic

Geochimico ef Cosmochimico Actn Vol. 43. pp. 861 lo 868 0 Pcrgamon Ras Ltd 1979. Prrnted in Grcar Britam

Gibbsite solubility and thermodynamic properties of hydroxy-aluminum ions in aqueous

solution at 25°C

HOWARD M. MAY, PHILIP A. HELMICE and MARION L. JACKSON

Department of Soil Science, University of Wisconsin, Madison, WI 53706, U.S.A.

(Received 1 May 1978; accepted in revised form 6 February 1979)

Abstract-Solubility curves were determined for a synthetic gibbsite and a natural gibbsite (Minas Gerais, Brazil) from pH 4 to 9, in 0.2% gibbsite suspensions in 0.01 M NaNO, that were buffered by low concentrations of non-complexing buffer agents. Equilibrium solubility was approached from oversaturation (in suspensions spiked with Al(NOs)s solution), and also from undersaturation in some synthetic gibbsite suspensions. Mononuclear Al ion concentrations and pH values were periodically determined. Within 1 month or less, data from over- and undersaturated suspensions of synthetic gibbsite converged to describe an equilibrium solubility curve. A downward shift of the solubility curve, beginning at pH 6.7, indicates that a phase more stable than gibbsite controls Al solubility in alkaline systems. Extrapolation of the initial portion of the high-pH side of the synthetic gibbsite solubility curve provides the first unified equilibrium experimental model of Al ion speciation in waters from pH 4 to 9.

The significant mononuclear ion species at equilibrium with gibbsite are Al’+, A10H2+, Al(OH); and AI(O and their ion activity products are *K,, = 1.29 x JO*, *K,, = 1.33 x 103, l K,, = 9.49 x JO-” and ‘@KS4 = 8.94 x 10-i’. The calculated standard Gibbs free energies of formation (AG;) for the synthetic gibbsite and the AIOH *+ Al(OH); and Al(OH); ions are -276.0, - 166.9, -216.5 , and -313.5 kcal mol- ‘, respectively. These AGF values are based on the recently revised AC,” value for A13+ (- 117.0 + 0.3 kcal mol-‘) and carry the same uncertainty. The AGF of the natural gibbsite is -275.1 f 0.4 kcalmol-‘, which suggests that a range of AG; values can exist even for relatively simple natural minerals.

INTRODUCTION

A QUANTITATNE understanding of the aqueous geo- chemistry of aluminous minerals and of the forms and reactions of aluminum ions in natural waters requires accurate thermodynamic constants describing the aqueous solubility and hydrolysis reactions of aluminum. Most previous sohtbility studies have yielded variable results for hydroxy-aluminum minerals because they have attempted to follow the kinetically slow and complex crystallization of Al(OH), precipitates produced by neutralization of Al salt solutions with base and because of difficulties in measuring concentrations of dissolved Al species in near-neutral solutions (HEM and ROBERSON, 1967; PARKS, 1972; SMITH and HEM, 1972).

Recent characterizations of important parameters governing crystallization of hydroxy-aluminum solids from such precipitates (TURNER and Ross 1970: Ross and TURNER, 1971) suggest that equilibrium solubility can be quickly reached in gibbsite suspensions slightly oversaturated with mononuclear dissolved Al. The equilibrium Al solution compositions thus attained can be used to determine thermodynamic properties of the suspended gibbsite and of the mononuclear Al ion species in solution. Achieving the same equilibrium solubility values from undersaturation will verify the determinations. The solubility determinations reported here are the first to precisely define the region of minimum solubility of gibbsite. It is this

861

portion of the solubility curve that defines the properties of the heretofore incompletely characterized Al(OH); and Al(OH); ion species.

Materials

EXPERIMENTAL

Two samples of gibbsite were used. One was a synthetic product (Baker and Adamson purified powder, lot No. C356Zll6J) and the other a natural mineral sample (Minas Gerais, Brazil; Wards Natural Science Establishment, Inc., Rochester, NY 14603). Samples of each were ground with a corundum mortar and pestle and wet sieved with distilled water through a 50pm brass screen. After sieving, the Minas Gerais sample was treated twice with citrate-bicar- bonate-dithionite solution to remove free iron oxides (JACKSON, 1975). Both gibbsite samples were then subjected to seven cycles of suspension in deionized water followed by centrifugation to remove particles smaller than 2pm and any soluble impurities (JACKSON, 1975). The resulting 2-50pm sized samples were then dried and stored for use.

Both of the treated gibbsites were analyzed for crystallinity and chemical purity. X-ray diffraction patterns for both materials exhibited only sharp, standard peaks for gibbsite, up to 40’ 20. The diffraction pattern for the synthetic mineral displayed somewhat more intense peaks and had a lower background, suggesting greater crystallinity for the synthetic material. Both gibbsites appeared as com- pact, euhedral crystals in scanning electron micrographs. Analysis of the samples for more than 25 elements by neutron activation (KOONS and HELMKE, 1978) showed only the following impurities: Baker and Adamson synthetic gibbsite-11 ppm Ga, 0.8 ppm Cr and 0.04 ppm Sb; Minas Gerais natural gibbsite-135 ppm U, 113 ppm Zn, 80 ppm Fe, 3.7ppm Cr. 2.9ppm As, 2.3ppm SC, 0.14ppm Sb and 0.05 ppm Co.

862 HOWARD M. MAY, PHILIP A. HELMKE and MARION L. JACKSON

The buffering agents selected to control the pH in the experimental gibbsite suspensions were: acetic acid (pK, = 4.76); 2.2-bis-(hydroxymethyl)-22, 2”nitrilotrieth- anal, ‘Bis-tris’, (pK, = 6.5): and (trishydroxymethyl)- aminomethane, ‘Tris’. (pK, = 8.06) (PERRIN and DEMPSEY, 1974). At the concentrations used in the solubility rtms. these buffer agents cannot significantly complex Al ion species (SILLEN and MARTELL, 1971; PERRIN and DEMPSEY, 1974). Ail buffering agents and the reagents HNO, and NaNO, did not contribute to or interfere with the Al analyses, with the exception of Bis-tris. All components of the experimental systems were rigorously tested in this respect. An impurity (probably Fe), which contributed to measured concentrations’ of A1 was removed from stock solutions of Bis-tris by extraction with 8-hydroxyquinoiine into toiuene, followed by successive toluene extractions to remove traces of 8-hydro~quinoiine from the Bii-tris stock solution. Sodium acetate and acetic acid used for the acetate buffer were_ ultrapure reagents (Aifa Products, Danvers, MA) and Bis-tris (99%) and Tris (99.9+x) were Gold Label high purity products of Aldrich Chemical Corp., Milwaukee, WI. Ultrapure distilled-deionized water was used for ail suspensioris and reagents. Toiuene used for solvent extraction and the HNOa used to construct Bis-tris and Tris buffers were both A.C.S. reagent grade.

Solubiiity dererminmions

Each soiubiiity determination was begun by suspending sufficient gibbsite to make a 0.2% suspension in solutions that were 0.01 M NaNO, and buffered at desired pH values with 0.001 M acetate, or 0.002M Bis-tris, or 0.0025 M Tris. In the systems buffered with acetate (pH 4.0-6.51, the gibbsite for each run wag initially given three, 6 hr washings with the same acetate-buffered NaNO, soiu- tion in which it would later be suspended for the soiubility run. These washings were intended to ‘titrate’ the gibbsite particle surfaces and thereby reduce pH drjft during equilibration. Subsequent determinations in the ranges pH 6-7 (Bis-trisf and pH 7-9 (Tris) attemRted to achieve the same objective by the use of higher buffer concentrations without the pre-equilibration washings. For solubility determinations from oversaturation, small aiiquots of Al(NOs)s solution were added to each buffered system at the same time as the gibbsite addition. The Ai(NOs)s aliquot addi- tions were calculated to set the initial dissolved Al concentrations about ten times higher than that predicted by the theoretical equ~ibr~~ gibbsite sofubiiity curve of SMITH and HEM (1972). Immediately after their preparation in polyethylene bottles, the suspensions were placed in a thermostatic sample shaker and agitated at 25.0 f 0.2%

Analysis

S&samples of the gibbsite suspensions were analyzed periodically from 2hr to 31 days after preparation. Depending upon the volume of solution aiiquot required for the determination of soluble aluminum, llOm1 suspension samples (in 125ml Teflon bottles) or 2Oml samples (in 45 ml polypropylene ‘Oak Ridge’ centrifuge tubes) were taken for each analysis. After pH measurement, the samples were centrifuged. The concentration of total dissolved mononuclear aluminum (Al,) was determined spec- trophotometricaliy in aiiquots of 1, 10, 100 ml of the super- natant by the procedure of MAY et 01. (1979). This procedure exploits kinetic aspects of 8-hydroxyquinoiine complexation and solvent extraction of mononuclear Al ions so that the total concentration of mononuclear Al ions can be determined in a sample with no contributions from poiynuclear Al ions or aluminous solids. The detection limit of the procedure is about 0.2ppb Al,. The but&r con~trations in the Al standards analyzed with each set of samples were equivalent to those in the sample aliquots to cancel any possible ipterference effects. Alkaline samples received small amounts of dilute HNOa immediately fol-

lowing the addition of 8-hy~oxyquinoline reagent, and before toiuene addition and extraction, to ensure that the analytical extraction of the AI-hydroxyquinoline comptex occurred at pH 5.0.

Suspension pH values were determined with an expanded scale meter and micro-combination electrode calibrated over narrow pH ranges (pH 4-7 or 7-9) with fresh standard buffer solutions. All calibrations and experimental pH det~inatio~ were made at 25.0 + 0.2% in a regulated water bath. The maximum uncertainty associ- ated with pH measurements is estimated to be +0.02 pH units.

RESULTS

The results for the solubility of the synthetic gibbsite from over- and under-saturation define a unique solubility curve as shown in Fig. 1. All of the initial and experimental solution composition points for the pH 4.0-U and pH 7-9 systems are shown as well as the experimental points for the pH 4-7 systems. Lines connecting the data points in the first two sets show the evolution of the solution composition for each individual run. As the solubility data were being collected for the pH 4.0-6.5 and 7-9 sets, an unex- pected discontinuity was noticed between pH 6 and 7. The apparent gap between pH 6 and 7, considered together with the somewhat longer times required to reach constant Al, concentrations in the alkaline suspensions, suggested that gibbsite is not a stable phase in alkaline systems.

To resolve the apparently discontinuous nature of gibbsite solubihty and aluminum solution chemistry in the range pH 6-7, oversaturation runs were set up at approximately 0.1 pH unit intervals between pH

Fig. 1. Soiubility data for synthetic gibbsite obtained from over- and undersaturation, as a function of pH and time. The dashed line representing gib~it~Al(OH)~ equilibrium is extrapolated from data below pH 6.7. The caicu- iated concentrations of the individual mononuclear Al ion species in equilibrium with gibbsite are shown as functions

of pH.

Hydroxy-aluminum ions in aqueous solution at 25°C 863

A 2hr

A 24 hr

A 192 hr

6 6.5 7

w

Fig. 2. Detail of data at the minimum solubility of gibbsite showing the rapid attainment of equilibrium and the inflection in the solubility curve.

6 and 7 and analyzed periodically. Both initial and experimental solution composition points are shown

in Fig. 2, together with the nearby solutions points from the previous two sets of runs. Again, all of the points for each run are connected by lines to depict evolution of the solution compositions. The curve thus described as a function of pH clearly has an inflection near pH 6.7 and begins to fall away from the initial trend begun between pH 6.0 and 6.7. A theoretical boundary representing the dissolution of gibbsite to aluminate ion [Al(OH);] can be drawn through the inflection at about pH 6.7, as shown by the dashed line in Figs. 1 and 2.

The chemical thermodynamic properties of the possible mononuclear Al ion species (A13+, A10H2*, Al(OH);, Al(OH)z and Al(OH);) coexisting at equilibrium with the solid phase can be determined from the gibbsite solubility curve shown in Fig. 1. The experimentally measured Al concentrations are the sums of all the mononuclear species present:

[Al-J = [A13+] + [AlOH*+] + [ANOH);]

+ CAW-%I + CWW;I (1) because contributions from polynuclear Al ion species and colloidal or solid Al forms were excluded (TURNER, 1969; MAY et al., 1979). As there is some doubt about the existence of an uncharged mononuclear Al species Al(OH): (P~RK$ 1972; SMITH and HEM, 1972) that species may be omitted in an altema- tive speciation model:

[Air] = [A13+] + [AlOH’+]

+ [AI(O + [ANOH);]. (2)

The models represented by eqns (1) and (2) were tested against the experimental data to evaluate the thermodynamic properties of the most probable par- ticipating ion species. This was done by relating each ion species to gibbsite through the appropriate ion

activity products *KSO, *KS,, *Ks2, *KS3 and *KS, for the reactions:

Al(OH)3 + 3H’ *Al”+ + 3HrO (3)

Al(OH), + 2H’ $ AlOH’+ + 2H20. (4)

Al(OH)3 e H+ F? Al(OH); + Hz0 (5)

AI it Al(OH); (6)

Al(OH)3 + H20 it Al(OH); + H’ (7)

and doing correlation analyses of the relationships between [AI,] and H’ activity. Thus, from eqn (1) :

*K,dH + J3 CAbI = yA13’ + yAyFH::

+K,z(H+) + ,,K

+ yAl(OH); s3 + *K,dH+)- ’

yAl(OH); (8)

where the four activity coefficients shown are 0.430. 0.664, 0.903 and 0.894, respectively, as calculated from the extended Debye-Hiickel expression (ROBINSON and STOKES. 1959). using the same ion size parameters a as did SMIIH and HEM (1972) and ionic strength I = 0.011 for all data. Equation (1) then becomes:

[Al,] = 2.33 *K,o(H+)3 + 1.506 *K,I(H+)2

+ 1.107 *Ks2(H+) + *KS3 + 1.119 *K,,(H+)- ‘. (9)

Equation (2) can be stated in similar form, differing only by omission of the Al( term, *KS3 :

[AIT] = 2.33 *K,0(H+)3 + 1.506 *K,,(H+)2

+ 1.107*Ks2(H+) + l.l19*K,*(H+)-‘. (10)

Equations (9) and (10) are polynomials which can be fitted to the experimental data by generalized least square procedures, The data employed for such com- putations are given in Table 1. The first ten paired values in Table 1 are the final (422 hr) solution compositions for the pH 4.0-6.5 acetate-buffered set of

864 HOWARD M. MAY, PHILIP A. HELMKE and MARION L. JACKSON

Table 1. Equilibrium and extrapolated solubility data used to calculate gibbsite ion activity products *K,, through *K,*

Equilibrium paired values? (H+) CAbI, M

Extrapolated paired values1 W+) CAbI, M

3.31 X 10-S 1.38 x lO-s 5.62 x lo-’ 2.55 x 1o-8 1.70 X 1o-5 2.15 x 1o-6 3.16 x lo-’ 3.62 x lo-* 1.02 X 10-s 6.35 x lo-’ 1.78 x lo-’ 5.92 x lo-’ 1.24 x lO-6 3.01 x lo-’ 1.00 x lo-’ 1.01 x lo-’ 4.57 x lo+ 1.24 x lo-’ 5.62 x lo-* 1.78 x lo-’ 2.34 x 1O-6 4.3 x 10-s 3.16 x lo-’ 3.22 x lo-’ 2.24 x lo+ 3.9 x 10-S 1.00 x lo-* 1.00 x lo-6 1.10 x lo-+ 2.4 x 1O-8 3.16 x lo-’ 3.22 x 1O-6 9.12 x lo-’ 2.4 x lo-* 1.00 x 1o-q 1.00 x lo-5 6.92 x lo-’ 2.4 x lo-’ - -

t Experimental data from acetate-buffered systems equilibrated 422 hr. $ Data based on extrapolation of projected theoretical gibbsite-Al(OH);

equilibrium phase boundary.

oversaturation runs, excluding the composition of the pH 4.0 system, which had clearly not reached equilibrium, as discussed later. The last 9 paired values in Table 1 are theoretical values, based upon the extrapolated position of the high-pH side of the gibbsite solubility curve (see Fig. 1).

Computer-assisted data analyses employed two linear least squares subroutines to attempt fits of the data in Table 1 to models representing eqns (9) and (10) above. The AIGILS subroutine uses generalized inverse concepts and AILLSQ uses Householder’s transformations (MACC Approximation and Interpo- lation Manual, 1972). Use of two different approaches serves as a check on the validity of the computed solutions, since equation coefficients computed by both approaches should be nearly identical. Values of y ([Al,]) were transformed l/y2 to partially offset the’effect of the wide range of y values on the sums

of squares calculations in the programs. The sets of polynomial coefficients computed with both subroutines are virtually identical, and are 2.95 x IO*, 2.23 x 103, 8.77 x 10-3, 2.23 x lO-9 and 9.93 x lo-l5 for eqn (9), and 3.15 x lOs, . 1.71 x lo”, 1.12 x low2 and 1.01 x IO-l4 for eqn (10). Consider- ation of the relative size of the *& term in eqn (9) and the near-equality of the remaining coefficients for the two equation models shows that Al( is an insignificant ion species. It never constitutes more than 10% of [Al,], if it exists at all.

The equation coefficients were also determined by simple numerical approximation because the precise effects of the weighting procedures used in the computer fit are difficult to assess, and because the computed coefficients generate solubility curves that deviate locally from experimental [Al,] values. The approximation is based on two assumptions. First, that the extrapolated portion of the gibbsite-Al(OH); line above pH 7 (broken line in Fig. 1) is valid and, second, that the relative error is approximately constant for all values of [Al,] along the solubility curve in Fig. 1. The second assumption recognizes that individual experimental [Al,] values carry analytical un-

certainties of +2% near pH 4 and f 17% near pH 6, but there are relatively more data points in the region of minimum solubility, which increases confidence in the accuracy of this portion of the solubility curve. The goal of this approximation was to generate coefficients for eqn (10) which fit the entire gibbsite solubility curve in Fig. 1, from pH 4.5 upward, as closely as possible. Assuming the value of the last coefficient to be 1.00 x lo- 14, the best fit of the eqn (10) model to the experimental solubility curve occurs when the first three coefficients are 3.0 x lo*, 2.0 x lo3 and 1.05 x IOm2, respectively. These equation coefficients lie between the two sets generated by the least squares subroutines for models with and without the Al( ion species. They produce a curve differing from the experimental curve by less than 3% at all points above pH 4.5. Thus, the equation:

[AIT] = 3.0 x 10“(H+)3 + 2.0 x 103(H+)2

+ 1.05 x 10-2(H+) + 1.00 x 10-14(H+)-L (11)

is taken to be the best fit of the experimental data and is the one used for subsequent calculations.

Thermodynamic properties of aqueous mononuclear aluminum ions

Multiplying the coefficients in eqn (11) by activity coefficients for the appropriate Al ion species yields the thermodynamic solubility products *&, to *K,4. From the appropriate *K, values, standard Gibbs free energies of reaction where then calculated for each of the above dissolutions, and combined with the standard Gibbs free energies of formation for the Al’+ ion, AG,” = - 116.97 f 0.33 kcal mol- ’ (HEMINGWAY and ROBE, 1977b), and for H,O, AC,” = -56.69 kcal mol-’ (ROBIE and WALDBAUM, 1968) to give AG,” values for the synthetic gibbsite and all of the Al ion species. These values, together with their maximum probable uncertainties of +0.3 kcal mol- ‘- which arises from the uncertainty in the value for A13+, are given in Table 2, along with the *K, values from


Table 2. Experimentally determined *K, and calculated AGF values for synthetic gibbsite and mononuclear hydroxy- aluminum ions

Al species

Ion activity product

AG;, kcal mol-’

Al(OH), gibbsite

*I& = 1.29 x 10s

-276 f 0.3

AlOH*+ Al(OH): Al(OH);

*?L., = 1.33 X 103 *I;.2 = 9.49 x 10-3 *I&, = 8.94 x lo-”

- 166.9 + 0.3 -216.5 + 0.3 -313.5 x 0.3

t Based on the values AGF = - 116.97 + 0.33 kcal mol- ’ for A13+ (HEMINGWAY and ROBE, 1977b) and AC; = -56.69 kcal mol-’ for Hz0 (ROBIE and WALDBAUM, 1968).

which they were calculated. Consideration of analyti- the synthetic mineral systems. The Minas Gerais cal uncertainties and the ranges of the coefficients for gibbsite is about four times more soluble than the eqn (11) indicates that introduced errors will be synthetic mineral. Analysis of the early and late sets 0.03 kcal mol- ’ or less for gibbsite and Al(OH); ion, of solubility data indicates that the value of AC; for and 0.1 kcal mol- ’ or less for AlOH’+ and Al(OH); the natural gibbsite ranges between -275.02 and ions. Therefore, the much larger 0.3 kcal mol- ’ error - 275.27 kcal mol- ‘. Considering uncertainties due in the AC,” value for A13+ is assumed to approximate to the effect of the pH shifts, the best value of AC: the maximum probable error in the calculated AC; for the Minas Gerais gibbsite is -275.1 + 0.4 k& values. mol-‘.

The initial and experimental solution composition points for the acetate-buffered and Tris-buffered suspensions of the Minas Gerais gibbsite are shown in Fig. 3. The solubility curve for the synthetic gibbsite is also given. Two slightly different solubility curves, each congruent with the synthetic gibbsite curve, are shown for both early and late sets of suspension analyses. As is evident from Fig. 3, the solution compositions in the acetate-buffered natural gibbsite suspensions drifted toward pH 7 and data from the pH 7-9 Tris-buffered suspensions also indicate pH drift with slow movement toward an equilibrium solubility curve. An apparent change in the solid phase controlling the concentrations of Air is discernible for the natural gibbsite above pH 6.7, just as observed in

The large shifts of suspension pH values argued against further detailed studies with the Minas Gerais gibbsite, but two useful points are demonstrated by these data. First. the congruency of solubility curves for both minerals reinforces the accuracy of the solubility experiments and, hence, the accuracy of the calculated molar free energies of the four mononuclear aluminum ions. Second, the solubilities of naturally occurring relatively pure minerals, such as gibbsite, may differ significantly from those of more nearly perfect synthetic specimens. Such variations in soiubility must be considered in geochemical modelling of ter- restrial environments.

Hydroxy-aluminum solids stable in alkaline systems

The solubility curves for both gibbsites suggest that a solid phase less soluble than gibbsite controls the

0 Initial q Initial

028hr -448 * 50hr hr

0

l 120 hr

384 24 2 hr hr hr

= 192hr

9 . P

4 5 6 7 8 9

PH

Fig. 3. Solubility data for a natural gibbsite as a function of pH and time. Solubility curves for early and late sets of suspension analyses (broken lines) are shown for comparison with the solubility curve

for the synthetic gibbsite (solid line).

a66 HOWARD M. MAY. PHILIP A. HELMKE and MARION t. JACKSON

solution composition above pH 6.7. Data from initially over- and undersaturated suspensions of synthetic g&site precisely describe an equilibria solid- Al(OH); boundary (Fig. 1) for this secondary phase (*Ks4 ranged from 2.24 x 10-l’ to 2.52 x lo-I’), but the identity of the phase could not be determined experimentally by X-ray diffraction or scanning electron microscopy. For the phase boundary shown (solid line in Fig. l), the solubifity product *&, is 2.38 x lo- ls. This is equivalent to a AG,” value of - 276.8 f 0.3 kcal mol- ’ for an Al(OHIS mineral or a AG,” of - 220.1 f 0.3 kcal mol- ’ for an AlO mineral.

DISCUSSION

Our *& value for the synthetic gibbsite is close to values determined by solubility measurements for well characterized crystalline gibbsites (K~TTRI~K, 1966; SING& 1974). Therefore, the AG,” for our synthetic gibbsite, which is based upon HEMINGWAY and ROB& (1977b) new AGfo value for A13+, is identical with the recently revised calorimetric AGF value for gibbsite reported by HEMINGWAY and ROBIE (1977a) and also with the value selected by PARKS (1972) after adjustment. All of the aluminum species AC: values discussed or compared in this report are based upon or have been adjusted to conform with a value of AG,” of - 116.97 + 0.33 kcal mol- ’ for Al3+ (HYING- WAY and ROBE, 1977a,b), which appears to be the most accurate estimate presently available for this fundamental quantity.

As noted earlier in the Results section, the pH 4 synthetic gibbsite system originating from oversaturation failed to reach equilibrium within 422 hr. Solu- bility products calculated from the data for this system and plotted as a function of (time)-* extrapolate to an ~uilibrium value close to that obtained from the first parameter in eqn (11). This slower progress toward equilibrium solubility has been noted in systems with high dissolved Al con~ntratio~ (S~fff~ and HEM, 1972) and may be the result of a pH-dependent, limiting step in the organization of gibbsite surfaces. As shown in this study, gibbsite suspensions poised near the solubility curve minimum can reach equilibrium solubility within relatively short periods of time when only small fluxes of dissolving or crys- tallizing Al ions are involved. The much ionger time requirements for Al solubility experiments, discussed by HEMINGWAY er al. (1978), are probably artifacts of the strongly acidic or alkaline conditions employed in the studies they reviewed.

The value of AC;,” for AlOH’+ calcuIated from our experimental results is identical with the value selected by PARKS (1972). Our thermodynamic pK, for the first deprotonation of AlfOH&+is thus 4.99, which compares well with the widely accepted values of 4.98 (SCHOFIELD and TAYLOR, 1954) and 5.02 (FRINK and PEECI-I, 19631.

Disagreements among reported estimates of AG: for the Al(OH)l ion result from a lack of data accu-

rately describing the minimum region of the solubility-pH curve. Our equilibrium analytical data de- scribes an unique solubility curve minimum consider- ably lower than that suggested by GAYER et al. (1958). RAUPACH (1963), DEZELIC et al. (1971) or SMITH and HEM (1972). Therefore, the AGF value we calculate for AI( is 2 kcalmol-* more positive than the value estimated by PARKS (1972) and HEMINGWAY et al. (1978). The sharply curved region of minimum solubility precludes any significant contributions by an Al(OH)z species. Reports of the existence of this species are based upon me~urements in non-equilibrated systems (DEZELIC et al., 1971), or by competitive ligand techniques (NAZARENKO and NEVSKAYA, 1969) which are inappropriate to the very low [Al,‘J values and short equilibration times characteristic of the minimum solubility region.

The relatively greater solubility of the natural gibbsite is likely due to its crystailinity being slightly less perfect than that of the synthetic material, as sug gested by data from X-ray diffraction, and the pres- ence of impurities as shown by neutron activation analysis.

Consideration of the solubitity data obtained above pH 6 (Fig. 1) suggests explanations for the confusion among the values of AG,” for the AI( ion discussed by PERKS (1972). The solubility change beginning at pH 6.7 has not been noted in previous reports of gibbsite solubility determinations, although the aluminum hydroxide synthesis literature contains numerous accounts of pfl-dependent formation of various crystal forms (cf. SCHOEN and ROBERSON, 1970). The change we have observed shows that most previous attempts to evaluate the AG,” for AI( are invalid because the identity and correct AC,0 of the controlling phase is needed to correctly determine the AG,” of the Al(OH& ion. The range of *KS, values reported for gibbsite and other hydroxyaluminum minerals dissolved in alkaline solutions (RAUPACH, 1963; Kn-iwc~, 1956; APE??, 1970; PARKS, 1972) pre- vents identification, by simple comparison, of the phase responsible for our value of *K,,,. Of available accounts, only that of GAYER d al. (1958) employed similar experimental conditions. Their quite uniform results from pH 9 to 12 yielded a l &, value of 2.98 x lo- i5, which is close to our value. The *KS4 value reported by biTRICK (1966), and employed by HEMINGWAY et al. (1978) to recalculate AG,” for Al(OH);, is essentially based upon extra~latio~ of two data points obtained from solubility experiments that did not attain equilibrium.

We recognize that the shape and placement of the solubility curve could be affected by reactions between Al ions and the buffers used to control pH (LIND and HEM, 197.5). Several points of evidence show that this is not the case in our experimental systems. The buffers used are very weakly competitive ligands and were present at low concentrations. &cause the pH range of the buffered Tris-suspensions brackets the pK, value for Tris, any complexation

of Al by Tris should be pH dependent. This is not minerals by HEMINGWAY et al. (1978). We tentatively the case as the equilibrium solubility values between identify the controlling solid in our alkaline suspen- pH 7 and 9 define a straight line with the slope sions as boehmite. This choice is based upon the expected for control by Al(OH); ion. In addition, the rapidity with which equilibrium was attained in our inflections beginning at pH 6.7 are displayed on solu- experimental alkaline suspensions. However, the bility curves for both minerals, yet the inflection in possibility that very small, neoformed diaspore par- the Minas Gerais curve developed in suspensions con- ticles controlled the chemistry of our alkaline suspen- taining acetate rather than the Bis-tris buffer used in sions, due to particle size effects on mineral solubility the synthetic mineral suspensions. The solubility (LANGMUIR and WHI’ITEMORE, 1971; SMI?H and HEM, curves for each mineral are continuous from one 1972) cannot be discounted. buffer system to the next, as well as being congruent. The shape of the solubility curve is fixed by the

X-ray diffraction (XRD) and scanning electron thermodynamic properties of the mononuclear Al ion microscopy were used to search for the newly-formed species. The position of the curve varies along its hydroxyaluminum phase indicated by the solubility vertical axis as a function solely of the stability of data. The original gibbsite was also examined and the equilibrium mineral phase. Consideration of the compared with the alkali-equilibrated solids. solubility data in Figs. 1 and 3 suggests that the Debye-Scherrer XRD patterns were made for the apparent stability of gibbsite below pH 6.7 may be smallest size fractions of the alkali-treated solids. a result of kinetic phenomena and that gibbsite may No evidence was found for a new, discrete crystalline be metastable in acid systems with respect to the less phase. Gibbsite was the only detectable phase in both soluble phase (e.g. boehmite). The complete solubility original’ and alkali-equilibrated materials. This is curve for the less soluble mineral would extend not surprising because consideration of Fig. 1 indi- through a minimum below that for gibbsite and on cates that the newly-formed solid responsible for the to a position below that of gibbsite in the acid range. lowered solubility would constitute 0.1% or less of Thus, the occurrence of boehmite or diaspore with the total hydroxy-aluminum solids present, in the gibbsite in carbonate terrains (KENNEDY, 1959) and absence of any significant recrystallization of the in- the predominance of gibbsite in acidic weathering en- itial gibbsite. vironments may result from kinetic limitations on the

Evidence from laboratory studies and field investi- crystallization of an AlO polymorph under gations suggests that the solid phase in equilibrium acidic conditions. with Al(OH); in our experimental alkaline systems is either boehmite or diaspore. Although reported

Acknowledgements-We wish to thank R. M. GARRELS and

measurements of the relative stabilities of bayerite, F. T. MACKENZIE for valuable discussions; S. W. BAILEY for advice on XRD analyses; H. M. HULL for assistance

nordstrandite and gibbsite are not in agreement (Rus- with computer analyses of the solubility data; R. L. KORO-

SELL et al., 1955; BARNHISEL and RICH, 1965; APPS, TEV for performing the neutron activation analyses; and

1970; SCHOEN and ROBERSON, 1970) and kinetic con- J. E. THRFSHER for assistance with scanning electron micro-

trols on crystallization obscure some thermodynamic scopy. This research was supported in part by the School

relationships, the rarity of naturally occurring bayer- of Natural Resources, College of Agricultural and Life Sciences, University of Wisconsin, Madison, WI 53706,

ite and nordstrandite strongly suggest that the alka- under project 1123. and in part by the National Science

line-stable phase is not likely to be one of the Foundation, EAR-7619783-Jackson.

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Gibbsite solubility and thermodynamic