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American Mineralogist, Volume76,pages445-457, 1991

Revisedyalues or the thermodynamicpropertiesof boehmite,AIO(OH),and relatedspeciesand phasesn the systemAI-H-OBnucnS. Hr'urNcw,l,v,Rrcn.cnoA. RonrnU.S. GeologicalSurvey, Reston,Virginia 22092,U.5.A.

JonN A. AppsLawrenceBerkeley -aboratory,Berkeley,California 94720, U.S.A.

AnsrnlcrHeat capacity measurements re reported for a well-characterizedboehmite that differsignificantly rom results reportedearlierby Shomateand Cook (1946)for a monohydrateof alumina. It is suggestedhat the earlier measurementswere made on a sample hat wasa mixture of phasesand that use of that heat-capacityand derived thermodynamic databe discontinued. The entropy of boehmite derived in this study is 37 19 + 0. 0 J/(mol .K) at 298.15K.A review of recent27Alsolution NMR data and other experimentshas shown that themethod of preparationof Al-bearing solutions can significantly affect he concentrationofmonomeric Al species n solution. Becausehe proceduresby which Al solution concen-trations are measured n solubility studiesdetermine the quantity of so-calledmonomericspeciespresent,apparent differences n calculatedgibbsitestability most likely arisefromdiferences n experimentalprocedure ather than from differencesn crystallinity, as oftensuggested. review ofpublished solubility data for gibbsitesuggestshat the best valuesthat can be currently estimated rom that data for the Gibbs free energiesof formation ofAl(OHh and Al3+are - 1305.0+ 1.3and -489.8 t 4.0 kJ/mol, respectivelyHemingwayet al., 1978).Basedon our value for the entropy and accepting he recommendedGibbs free energyfor AI(OH);, we have calculated the Gibbs free energy and enthalpy of formation ofboehmite o be -918.4 + 2.1 and -996.4 + 2.2kl/mol respectively,rom solubilitydatafor boehmite. TheGibbs energy or boehmite is unchanged rom that givenby Hemingwayet al. (1978).

fNrnonucrroN monly attributed to boehmite (Shomateand Cook, 1946)Minerals with compositions in the chemical system were basedon heat-capacitydata for a phasedescribedAlrO3-SiOr-HrO are ft"settt in a wide variety of sedi- as a_monohydrate that producedan X-ray pattern similarmentary and metamorphic rocks and in soils. Metasta- to that of bayerite' Kelley and King (1961)were the firstbility of phases n this system has resultedin disparate to describe the phaseas boehmite. Subsequent abula-experimental data and consquentdisparate interpreta- tions of thermochemical data have followed Kelley andtion of the phaserelationshipsof some of the phaies in King ( I 96 )' Recently, Apps et al. ( I 988) concluded hatthe system.Boehmite,AIO(OH), is an exampleof suchu the entropJ @I 298.15 K and I bar) of boehmite was toophase. largeand Berman et al. (1985)concluded he heat capac-Diaspore,AIO(OH), is generallyacceptedo be the sta- ities reported for boehmite were oo large.We report hereble aluminum hydroxide or oxyhydroxide phaseat earth heat-capacity measurements for a well-characterizedsurface emperatures e.g.,Parks, 1972; perkins et al., sampleofboehmitethatvalidatetheconcernsexpressed1979;Hemingway,1982).However,conflictingconclu- by Parks 1972),Hemingwayet al. (1978),and Bermansions have beenreached egarding he relative stabilities et al. (1985) and confirm the prediction of Apps et al.of other aluminum hydroxides and oxyhydroxides (e.g., (1988).Sanjuanand Michard, 1987,have calculateda Gibbs en-ergy for bayerite that would make it slightly more stable Samplethan diaspore). Inconsistenciesand errors in the ther- Two samplesof boehmite weresynthesizedby K. We-modynamic data utilized by the severalgroups may be fersofAlcoa. The first samplewasprepared rom gibbsite

responsible or someof the disparate nterpretations. Parks during a 20-h hydrothermal experient at approximately(1972) and Hemingway et al. (1978)noted and expressed 473 K. The resulting boehmite was well crystallized, asconcern that the entropy and heat-capacitydata com- shown by X-ray diffraction and SEM analysis.Aggregates0003-004X/9/030,1-0445$02.00 445

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446 HEMINGWAY ET AL.: THERMODYNAMIC PROPERTIES OF BOEHMITETABLE . Chemical nalysis or syntheticboehmite

Abundance'o/"

Trau 2. Experimental ow-temperature molar heat capacitiesfor boehmiteTemper- Heat Temper- Heat Temper-ature capacity ature capacity atureK J/(molK) K J/(mol.K) KI

MgNaKAuBaCaCeCoCrCsDyFeGaHfLaMgMnMoNiRbScSiSmTaUYbZn

HeatcapacityJ/(mol.K)

ppm

- NA A analyses y HelenV. Michel.-'lCP analysesby Andrew W. Yee.

of crystalsaverage 5 pm in diameterwith the individualcrystallites0.2-2 pm in size and displaying well-definedcrystal outlines.No evidenceof amorphous material wasfound; however, two additional phases [diaspore,AIO(OH), and akdalaite,4Al,O3 H,Ol were ound n smallquantity. This sample was used or all calorimetric stud-ies reported here.Following identification of diaspore n the first sample,a secondsample was synthesized rom the same startingmaterial, at the same temperature,but for a shorter pe-riod of time. The resulting sample was well-crystallizedboehmite in the form of aggregates f very thin plateletsof individual crystallites.No evidenceof amorphous ma-terial or other phaseswas ound. However, the very smallcrystallites made this sample mpractical as a sample orcalorimetry becauseof the low effectivepacking densityand the possibility that the He exchangegas would beadsorbedon the fine crystallitesat temperatures ess hanapproximately 15 K.A chemicalanalysisof the boehmitesample s listed inTable l. Combining the results of SEM, X-ray, andchemical analyses,we estimate that the sample contains< l0locombined diaspore, AIO(OH), and akdalaite,4Alr03.HrO.Portions of the first sample were used for both low-temperatureand differential scanningcalorimetric (DSC)measurements. he sampleused or low temperatureheat-capacitymeasurements as 27.3540g. Superambientheat

43.4(22146.2(23)*o.12('t2)0.814(8)

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HEMINGWAY ET AL.: THERMODYNAMIC PROPERTIES OF BOEHMITE 447Hemingway, 1972),bu;tnot for the heat capacitycontri-butions from the diaspore and akdalaite contaminantphases hat are discussedbelow. The observed heat ca-pacitiesweresmoothedusing cubic splinesmoothingrou-tines and weregraphicallyextrapolatedas Cr/T vs. Z2 to0 K using he experimentaland smoothedvalues or tem-peraturesbelow 30 K. Smoothed values of the heat ca-pacitiesand derived thermodynamic functions are listedin Table 3.The heat capacities of diaspore (Perkins et al., 1979)are smaller than those of the impure boehmite sample atall temperaturesbelow 300 K. Therefore,corrections tothe observedheatcapacitiesof the impure boehmite sam-ple for the contribution of diasporewould result in anincrease n the calculated heat capacitiesand entropies.A correction of approximately +0.02 J/(mol.K) wouldbe required in the entropy of boehmite at 298.15 K foreach loloof diaspore.Low-temperature heat capacities have not been re-ported for akdalaite;however,Mukaibo et al. (1969)havereported heat-capacitymeasurements or tohdite at su-perambient temperatures hat may be used to estimatethe magnitude of the correction for the akdalaite impu-rity. (Note that akdalaite and tohdite are considered obe the samephase,e.g.,Fleischer, 197 .) The specificheatof akdalaite is approximately 4o/o reater han that of co-rundum at 298.15 K. This difference s expected o in-creaseat lower temperatures.Basedon a comparison withboehmite and diaspore, we estimate the averagediffer-ence n the specificheatofakdalaite between0 and 298. 5K to be l0o/0.Using the entropy of corundum from Robieetal. (1979),we estimate he entropy of akdalaite 4AlrO3.H,O) to be 234 J/(mol.K) at 298.15 K. Basedon thisestimatefor the entropy of akdalaite, we estimate that acorrection of +0.04 J/(mol.K) would be required in theentropyof boehmiteat298.15 K for each 0lo f akdalaitermpunty.We estimate the total correction to the entropy calcu-lated for the impure boehmite sample a|298.15 K (Table3) to be less han 0.06 J/(mol.K). Becausehe boehmitesample contained a small quantity of very small crystalsthat would havea slightly higher specificheat than coars-er material of the same composition and becausewe es-timate the uncertainty n the calculatedboehmite entropyat 298.15K to be +0.I J/(mol.K), we have not applieda correction for the observed mpurity phases diasporeand akdalaite).The measurementsreported by Shomate and Cook(1946) for a monohydrate of Al and variously attributedto the phaseboehmite e.9.,Kelleyand King, l96l) maybe comparedwith the heat capacities eported herein. Atall temperatures opresentedby their measurements,heresults of Shomateand Cook (1946)are significantly arg-er than those reported in Table 2. At 100 K the heatcapacitiesdiffer by approximately 44o/o nd by approxi-mately 2lo/oat 298.15K. Shomateand Cook (1946)alsoreported heatcapacitydata for gibbsite hat may be com-pared with the data for gibbsite reported by Hemingway

TABLe . Molar thermodynamicpropertiesof boehmite

Heatcapacitycg EntropyS"r 58

GibbsEnthalpy energyfunction function(t/f, -(G9 -Hlllr r$lrTemperatureK J/(mol.K)51 01520253035404550607080901001 1 012013 014 015 016 0170180190200210220230240250260270280290300310320330340273.15298.15

0.0030.0270.1 30.1760.2640.4400.6961.0051.4441.8973.0524.5156.2748.31110.5812.9815.4818.O220.5923.1425.6828.1930.6833.1335.4937.7339.8841.9543.9545.8847.7349.5251.2552.9354.53s6.0657.5158.9360.3350.0754.24

0.0010.0080.0410.0880.1350.197o.2840.3950.5390.7141.1571-7332.4483.3014.2935.4136.6497.9889.41810.9312.5014.1315.8117.5419.3021.0922.8924.7126.5428.3730.2132.O433.8735.7037.5239.3441.1442.9344.7132.6237.19

0.001 0.0000.006 0.0020.033 0.0080.065 0.0230.094 0.0410.136 0.0610.198 0.0870.277 0.1180.382 0.1560.511 0.2030.833 0.3231.251 0.4821.766 0.6822.377 0.9243.082 1 2103.872 1.5404.735 1.9145.659 2.3296.634 2.7847.649 3.2768.697 3.8039.770 4.36210.86 4.95211.97 5.56813.09 6.21114.21 6.87715.33 7.56316.44 8.26917.54 8.99218.64 9.73119.72 10.4820.79 11.2521.85 12.02

22.89 12.8123.92 13.6024.93 14.4025.93 15.2126.91 16.0227.87 16.8421.13 11 4923.73 13.45Note-'Molar mass : 59.989g

et al. (1971),who used a calorimeter that was similar tothe oneused n this study.The data of Shomateand Cook(1946)were higherat all temperatures,but the maximumdifferencewas approximately 7o/o I 52.8K and that dif-ference decreased o l.5olo at 298.15 K. The foregoingcomparison suggestshat the differences etween he heatcapacity values reported by Shomate and Cook (1946)and the values reported n this study ariseprimarily fromdifferences n the sample, not in the equipment or pro-cedures or data processing.

ffrcrr-rnupnRATuRE HEAT cApAcrrrEs ANDTHERMODYNAMIC FUNCTIONS

Experimentalsuperambientheat capacities or boehm-ite are listed in Table 4. The results representmeasure-ments based on several samples,with new samplespre-pared and used following any partial dehydration of asample. Although the majority of HrO was lost in the

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448 HEMINGWAY ET AL.: THERMODYNAMIC PROPERTIES OF BOEHMITE

Series1338.9 59.77348.9 60.94358.8 61.93368.8 63.02378.7 64.05388.6 65.00398.6 66.02408.5 66.99418.5 67.95428.4 68.85438.3 69.81448.3 70.48458.2 71.50468.1 72.24478.1 73.03488.0 73.90497.0 74.66

Series2338.9 59.82348.9 60.97358.8 62.03368.8 63.08378.7 64.09388.6 65.07398.6 66.07

Series 2408.5 67.07418.5 67.97428.4 68.91438.3 69.83448.3 70.57458.2 71.61468.1 72.37478.1 73.18488.0 74.00497.0 74.85Series 3607.3 79.86616.2 81.35Series 4468.1 72. ' t1478.1 72.75Series 5597.3 79.53607.3 79.88616.2 81 79Series 6537.7 77.39547 6 77.76557.6 78.24567.5 78.53

Selies6577.5 78.93587.4 79.05597.3 79.40607.3 79.81617.2 80.31627.2 80.91637.1 81 14647.0 81 73657.0 82.36666.9 83.21Series7537.7 77.40547.6 77.77557.6 78.20567.5 78.63577.5 78.97587.4 79.32597.3 79.61607.3 80.30617.2 80.55627.2 81.11637.1 81.44647.0 81.69657.0 82.56666.9 83.54

TABLE, Experimentaluperambient olarheatcapacitiesorboehmiteTem- Heat Tem- Heat Tem- Heatperature capacity perature capacity perature capacityK J/(mol.K) K J/(mol.K) K J/(mot.K)

Between300 and approximately 420 K, the heat capac-ities derived from the data of Shomate and Cook (1946)are arger han those eported here.Above approximately420 K. they are smaller.As noted earlier, the most reasonableexplanation forthe difference between the data presentedby Shomateand Cook (19a6) for a monohydrateof Al and the resultspresentedhere is that the monohydrate studied by Sho-mate and Cook (1946) was not pure boehmite. Bayeritewas dentified by X-ray analysis o be a constituentof thesample studied by Shomate and Cook (1946). Mukaiboet al. ( I 969) g1ve 73 K as he temperatureof dehydrationof bayerite, consistentwith the observation of Shomateand Cook (1946) thar significant HrO loss occurred attemperaturesat or below 520 K. In addition, the densityof the sample used by Shomateand Cook (1946) wasdetermined and found to be 2.83 g/cm3, considerablylower than the theoreticalvalue of 3.07g/cm3 or boehm-ite; however, he value of 2.45g/cm3 eported n the samestudy for gibbsite is identical with the theoretical valuefor gibbsite, suggesting hat analytical error was not afactor.The procedurereported by Shomateand Cook (1946)for the preparationof their monohydrate alumina samplewas ollowed as closely aspossible.Dehydration ofgibbs-ite yielded a fine-grained,poorly crystallizedmixture ofboehmite and akdalaite.The presence f akdalaite n oursynthesismay indicate a possiblecauseof the nonstoichi-ometry (low HrO content) of the samplepreparedby Sho-mate and Cook (1946).Also, the addition of H,O to oursample with subsequentheating at 80 "C (following theprocedure of Shomate and Cook, 1946) yielded strongX-ray peaks or bayerite and significant oss of boehmiteX-ray peak intensities. Thus, several lines of evidencesuggesthat the monohydrate samplestudied by Shomateand Cook (1946) was a mixture of phases.

Fnrn nNpncv oF BoEHMTTEThe entropy derived in this study cannot be combinedwith the enthalpy of formation of boehmite to estimatedirectly the Gibbs energyof formation, as the enthalpyof formation of boehmite has not been established nde-pendently.Estimates of the Gibbs freeenergyof boehm-

ite may be developed rom phaseequilibria and solubilitydata.Ervin and Osborn (1951) developeda phasediagramfor the system AlrO3-HrO based upon a study of unre-versedsynthesisexperiments.Fields of stability are shownfor gibbsite,boehmite, diaspore,and corundum basedonthe products of crystallization of AlrO, gel and 7-alumi-nia. Equilibrium was considered to be proved if bothstarting materials (i.e., structurally different materials)yielded the sameproduct. Boehmite commonly formedand then slowly recrystallized o form diasporeor corun-dum in the P-Iregions designatedas he stability regionsfor those phases.Therefore, the synthesis experimentsprovide only a limit as to the minimum or maximumfree energy hat a phasemay have. Realistically, the ex-

Note.'Molarmass 59.989g.

temperature interval of 680-710 K, some loss occurredas ow as 450-600 K and resulted n the calculatedheatcapacities or that interval having a U-shapedcurvature.Weight loss calculatedafter such experimentswas gen-erally ofthe order ofa few hundredthsofa percentofthesampleweight and probably representedossof adsorbedHrO from some of the very small crystals. Each scanpresented n Table 4 representsone continuous set ofmeasurementsor the averageof severalcontinuous setsof measurements.Smoothed values of the heat capacitiesand thermo-dynamic functions are isted in Table 5. Smoothedvaluesof the low-temperature heat capacitiesbetween 298.15and 350 K werecombined with the high-temperature eatcapacitiesand the combined data set was fit by leastsquares o a 4-term polynomial with the constraint thatthe equation exactly fit the smoothed heat capacity at298.15K. The resultantEquation l,

C":205.721 - 0.0349217" 2635.27T 5+ 1.02666 106Z ( l )fits the data with an averagedeviation of +0.30/oand isvalid from 298 to 600 K.The measurementspresentedhere may be comparedwith the data of Shomateand Cook (1946)who measuredthe heatcontenrof their sample rom 321 to 520 K. Sho-mate and Cook (1946)terminated their measurements tapproximately 520 K becauseHrO evolved irreversiblyfrom the sampleand condensed n the sample capsule.

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HEMINGWAY ET AL.: THERMODYNAMIC PROPERTIES OF BOEHMITETABLE . Molar thermodynamicproperties or boehmiteAIO(OH): rystds298.15o 600 K.

449

Heat capacity Formation rom elementsTemperature q .+ tf, - l+*"YT -(@, - l4""llr Enthatpy Gibbs reeenergy298.1

Gibbs energyEntropy Enthalpy unction function

300350400450500550600

54.24+0.1054.5161.0266.4170.8574.5177.5480.04r0.32

37.190+0.10037.52646.43354.94463.03070.69177.93984.797+0.320

0.0000.3358.55315.4592'1.37526.51131 01 735.001

37.19037.1 137.88039.48541.65644.18046.92249.796

-996.389+2.100-996.415-996.846-997.269-997.367-997.365-997.058-996.990+2.300

-918.400+2.200-917.916-904.720-891.595-878.351-865.153-851.917-838.740+2.400ft"" - l4 7.07s+ o.o2o JTransitionsnphase Molar volume 1.9535 0.001 /barTransitions n referencestate elementsAr M.P. 933.45KEquationsq:205.721 - 0.034921- 2635.277-051026660r+(Validange: 98.15 o 600 K; Average bsolute ercent eviation: .50)Note.'Molarmass 59.989g.

perimentally determinedphaseboundariesmay representonly those regions n which the kinetics of recrystalliza-tion are rapid enough o be observedduring the durationof the experimentsperformed by Ervin and Osborn (I 9 5 ),and boehmite may not be stable under any of the P-7conditionsstudiedby Ervin and Osborn(1951).We may use he experimental reactionboundariesgiv-en by Ervin and Osborn (1951) to calculateminimumand maximum free energy values for boehmite. Trans-formation of gibbsite to boehmite is estimatedto occurat approximately 400 K and 3 bars. Thus the minimumfreeenergyof boehmiteat 400 K is estimated o be - 883.0kJlmol at 400 K and -909.8 kJ/mol at 298.15K (usingancillary data from Robie et al., 1979).Transformationof boehmite to corundum is estimated to occur at ap-proximately 658 K and 136bars. From these esultsandancillary data from Robie et al. (1979), the maximumfree energy for boehmite is estimatedto be -922.1 kJ/mol at 298.15K. Similarly, the maximum freeenergy ordiaspore s estimatedto be -922.7 kJ/mol from the re-action boundary for the reaction 2 diaspore corundum+ HrO given by Ervin and Osborn (1951).Within theforegoingcalculation and in subsequentcalculations n-volving gibbsite and corundum, the thermodynamicproperties isted in Robie et al. (1979\ are assumed o begood estimatesof the true values for thesephasesand,therefore,are used as fixed values. Resultsand evalua-tions presented y Haaset al. (1981),Hemingwayet al.(1978),Hemingway (1982),and Apps er al. (1988) pro-vide support for this assumption.Boehmitewasconsideredby Kittrick (1969) o be morestable than gibbsite, basedon a comparison of thermo-chemical data (seealso Parks, 1972; and,Hemingway etal., 1978).However, Kittrick (1969) selected ree energyof formation valuesfor gibbsite and boehmite that were

basedon different values or the freeenergyofformationof Al3+. Correcting the data for gibbsite to the same Alreferencevalue internally consistentwith the boehmitedata set reverses he relative stability ofthe two phasesat 298.15K (assuming he activity of HrO is unity), butthe difference s less than the experimental uncertainty.The contrary conclusionofChesworth (1972), hat gibbs-ite and HrO at unit activity are stable with respecttoboehmite at near-surfaceconditions, must also be con-sideredsuspect ecauset is basedon the highly uncertainestimate or the Gibbs energyof boehmite given by Ros-sini et al. (1952).Parks 1972)and later Hemingwayet al. (1978)select-ed the solubility data of Russellet al. (1955) or boehmiteas the best data set from which to determine the freeenergyof formation of boehmite. The result of the cal-culation is - 9 8.4 + 2. kJ/mol for the Gibbs freeenergyof formation of boehmite at 298.15 K, and it suggeststhat boehmite * H,O is more stable han gibbsite. n thiscase, he assumptions made are (l) that the solubilityproduct at 298.15 K can be calculated from the extrap-olation of solubility data at higher temperatures, 2) thatthe solution speciesnvolved in the higher temperaturesolubility experiments s the same speciesas that com-monly referencedat lower temperatures,and (3) that thefree energyofformation is accepted or that species.The estimate of the Gibbs free energyof formation ofboehmite given by Hemingway et al. (1978) s consistentwith two recentstudies. Hovey et al. (1988) calculatedarevised value for the Gibbs free energyof formation ofAI(OH); from solubility data for boehmite and a valuefor the Gibbs free energy of formation of boehmite ob-tained from Apps et al. (1988, then in preparation).Theresult,ArGln, : -1305.6 kJ/mol, is in agreementwiththe value derived by Hemingway et al. (1978) but was

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450 HEMINGWAY ET AL.: THERMODYNAMIC PROPERTIES OF BOEHMITEderived somewhatcircularly.Apps et al. (1988)deriveda value for the Gibbs freeenergyof formation of boehm-ite from the thermodynamic propertiesof gibbsite(Hem-ingway and Robie, 1977a)and from solubility data forgibbsiteand boehmite. Apps et al. (1988)reported valuesof ArGln,of -917.5 and - 1304.8kJ/mol, respectively,for boehmite and AI(OH);. From theseresults, we mayonly conclude that a Gibbs free energy of formation forboehmite of - 918 kJ/mol is consistentwith the solubilitydata for gibbsiteand boehmite and with a valueof - 1305kJlmol for the Gibbs free energy of formation ofAI(OH);. However, since the value for the Gibbs freeenergyof boehmitedeterminedby Apps et al. (1988) sreferenced to solubility data interpreted to representgibbsitesolubility, the resultsand interpretations are sub-ject to the samequestions posed or the studiesof Hem-ingway et al. (1978)and Hemingway 1982) hat are dis-cussed n the next section.

DrscussroNRecent work by May et al. (1979), Couturier et al.(l 984),and Sanjuan nd Michard (l 987)havequestionedthe free energyof formation of AI(OH); derivedby Hem-ingwayet al. (1978)and Hemingway 1982).Specificallyat issue s the questionofthe phaseor phaseshat controlthe solubility of Al in solutions with pH > 6 at 298.15K and at higher temperatures.Of more generalconcernis the questionof the relative stabilities of the aluminumhydroxide and oxyhydroxide phasesand the mechanismsby which precipitation ofthese phasesare controlled. Inthe discussion hat follows, we provide a detailed reviewof research hat establishes he relative stability of theAI(OH)3 polymorphs (bayerite,nordstrandite,and gibbs-ite). We select a set of solubility data to representcon-ditions of metastableequilibrium betweengibbsite andAl solution speciesand calculate he Gibbs free energiesof formation of Al3+ and AI(OH);. Finally, we providean explanation for the apparentvariation ofgibbsite sol-ubility that has variously been attributed to grain size,acid pretreatment,or sample crystallinity.

Historical perspectiveHemingway and Robie (1977a) identified an error in

the calorimetric procedureupon which the enthalpy andfree energy of formation of gibbsite were based and re-ported a revised set of thermodynamic values for gibbs-ite. In subsequentwork, Hemingway and Robie (1977b)and Hemingway et al. (1978) reported revisedvaluesforthe free energy of formation of AF* and AI(OH); basedon solubility studies for gibbsite by Kittrick (1966) andSingh (1974) and reported the revised value for the freeenergyof gibbsite.Implicit in the work of Hemingway etal . (1978) s the assumption hat gibbsitecontrolled heAl solubility observedby Kittrick (1966).The validity ofthis assumption was first questioned n the study of Mayet al. (1979).May et al. (1979) determined solubilities for a naturaland a syntheticgibbsiteat severalpH valuesbetweenpH

4 and9 usingseveralorganicpH buffers.May et al. (1979)obtained two subparallel curves, one for each sample,that displayed offsets oward lower solubility at approx-imatelypH:7 . May et al. 1979)concludedhat, n basicsolutions, the solubility of Al was controlled by a phasemore stable than gibbsite (when synthetic gibbsite wasused as the starting material). May et al. (1979) tenta-tively identified the phaseas boehmite, although no ev-idence for a phaseother than gibbsitewas found. May etal. (1979)utilized the solubilities determined from mea-surements obtained from the natural gibbsite sample insolutionswith pH > 7 to calculatea revisedvalue for thefree energyof formation of AI(OH); and, subsequently,to calculate the free energy of boehmite as -920.9 kJ/mol. May et al. (1979) concluded that the difference nsolubilities observed between the natural and syntheticgibbsites n acid solutionswas a consequence fa differ-ence in crystallinity of the two samples;however, theyalso argued hat, in basic solutions,the similar differencein observedsolubilitiesresulted rom control of solubilityby two phases.The apparent nconsistenciesn interpre-tation of the solubility data led Hemingway (1982) toquestion he interpretations.Hemingway (1982) combined the resultsof May et al.(1979)with other solubility studies rom the literature toprovide an alternative explanation for the features ob-served n the solubility data of May et al . (1979).As Alhydrolysis progresses,he nature and characteristicsofthe aqueousAl specieschange.The observed offset insolubility curves correspondswith the change rom thedominanceof species raditionally considered o be of theform Al(OH)t3-v)+ o the form AI(OH); (e.g', BaesandMesmer, l98l). Hemingway concludedthat a change nthe mechanism of precipitation accompanied he changein species tructure.Hemingway(1982)reasoned hat thischange n precipitation mechanism allowed bayerite toprecipitate only in solutions with a pH greater han ap-proximately 6. Hemingway (1982) concluded hat in theexperimentsperformed by May et al. (1979) supersatu-ration with respect o phasesother than gibbsiteoccurredand that the subparallel solubility curyes resulted fromcontrol of solution concentration of Al by precipitationof AI(OH), phases,nordstrandite and gibbsite n the acidregion, and bayeriteand nordstrandite n the basicregion.Hemingway (1982) further concluded that bayerite wasthe least stable of the three A(OH)3 phasesand thatgibbsitewas the most stable.Sanjuan and Michard (1987) measured he solubilityof gibbsiteat 323 K using proceduressimilar to those ofMay et al. (1979), obtained a similar offset in the ob-served solubility curve, and concluded that both the in-terpretations of May et al. (1979)and Hemingway (1982)were incorrect. Their interpretation was based on evi-dencethat bayerite is either found in alkaline solutionsor replacesgibbsite in alkaline solutions (Verdesand Gout,1987;Schoen nd Roberson,1970) nd hat AI(OH); hasa stability constant similar to that reportedby May et al.(1979) (Couturieret al., 1984).Although both Sanjuan

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HEMINGWAY ET AL.: THERMODYNAMIC PROPERTIES OF BOEHMITE 451and Michard (1987)and Hemingway 1982) questionedthe phaseor phasescontrolling Al solubility in the alka-line region, they reacheddifferent conclusions with re-spect o the relative stability ofbayerite and gibbsite,andwith respect o the free energyassignedo AI(OH);. Thelatter difference, n particular, must be resolved beforethe free energyof boehmite can be calculated.Bis-Tris,an organicpH bufferusedby May et al. (1979)in their solubility studiesofgibbsite, hasa strong enden-cy to form a complex with the aluminate on (Wesolowskiet al., 1990).The increase n total dissolvedAl resultingfrom this process s thought by Wesolowskiet al. (1990)to be the causeof the offset observed in the solubilitycuryespublishedby May et al. (1979).Tris, the organicpH buffer usedby May er al. (1979)at high pH, does notshow a strong tendencyto complex aluminate ion (We-solowskiet al., 1990). Thus, the resultsof May et al.(1979) may not representequilibrium betweengibbsiteand only hydrolized Al solution speciesn the pH rangebuffered by Bis-Tris, and the presenceof a phasemorestable than gibbsite may not be needed to explain thesolubility curve for gibbsitepublishedby May et al.(1979)and by Sanjuanand Michard (1987).The enhanced ol-ubility of Al is consistentwith the argumentsof Heming-way (1982) and with the observation of bayerite in suchsolutionsby Verdesand Gout (1987).Solubilities ower than those ound by May et al.(1979),Singh 1974),and Kittrick (1966)were reportedby Bloomand Weaver(1982)and Peryeaand Kittrick (1988) forseveralgibbsitesamples hat had previouslybeen studiedby others(Kittrick, I 966; Frink and Peech,1962).Bloomand Weaver (1982) attributed the lower solubility theyobserved n acid solutions to the removal of fine crystalsof gibbsite or reactive surfacesby acid pretreatment ofeach sample, and they ascribed the downward shift ofsolubility seen n the results of May et al. (1979)to morerapid Ostwald ripening of gibbsite n basicsolutions.Bloom and Weaver(1982)have compared heir solu-bility resultswith one of several sets of results reportedby Frink and Peech(1962) also for acid pretreatedma-terials.Bloom and Weaver(1982)did, in fact, find a low-er solubility than the set they chose,but other data pro-vided by Frink and Peech(1962) are equivalent to thesolubility reported by Bloom and Weaver (1982). Hem-ingway (1982) compared the data of Frink and Peech(1962) to the model he proposed(his Fig. 5). AlCl, so-lutions to which Frink and Peech 1962) addedHCI andgibbsite,and that were aged one month, showedsolubil-ities equivalent to those reported by Bloom and Weaver(1982).AlCl, solutions to which no HCI or gibbsitewereaddedhydrolyzedand showedsolubilitiesafter 3 monthsthat were equivalent to those found by May et al. (1979)for their natural gibbsite. A similar solution to whichgibbsite was added showed solubilities for pH < 4 thatwere equivalent to the solubilities for the solution equil-ibrated without gibbsite,whereas hose with solution pH> 4 showedsolubilities equivalent to thosegiven by Mayet al. (1979) for syntheticgibbsite.Finally, solutions that

contained no added AlCl., but contained gibbsite andwere acidified with HCl, showed the lowest solubilities.The results reported by Bloom and Weaver (1982) areequivalent to the results reported by Frink and Peech(1962) where the same experimental parameters weremaintained.Bloom and Weaver (1982) observed a significant dif-ference n the solubility of two sized fractions of FisherACS A(OH).. The sample FC with the smaller size rac-tion showed the greatersolubility. The sample FF withthe larger size fraction was pretreated with dilute acidwhereassampleFC was not. The solubility differenceob-serVedby Bloom and Weaver (1982)was ascribed o theacid pretreatment processes.However, the solution inwhich sampleFC was suspendedwas 0.01 M KNO. andresulted n a somewhatdifferent chemistry for the studiesof FC and FF. The effect of this difference s discussedbelow.Bloom and Weaver(1982)have shown that acid pre-treatedgibbsitesamplesFF, C-730, and C-33 yield iden-tical solubility products. SamplesC-730 and C-33 werestudiedpreviouslyby Kittrick (1966,solubility study) andHemingway et al. (1978, solution calorimetry). Heming-way et al. (1978) also measured he enthalpy of solutionof Fisher A(OH)3 similar to samplesFC and FF. Bloomand Weaver (1982) and Hemingwayet al. (1978) con-cluded from their studies hat all of the gibbsite sampleshad equivalent free eneryies.Peryeaand Kittrick (1988) have used a similar proce-dure to that used by Bloom and Weaver (1982) to studythe solubility of corundum, gibbsite,boehmite, and dia-spore. Peryea and Kittrick (1988) found Al concentra-tions in apparentequilibrium with gibbsitesampleC-730lower than that reportedby Kittrick (1966),and in agree-ment with the results reported by Bloom and Weaver(1982).Peryea nd Kittrick (1988)calculated he freeen-ergy of formation of the four Al-bearing phasesbasedonthe solubilities they measuredusing the value of the freeenergyof Al3+ given by Hemingway et al. (1978).Theresultsof these calculations were free energyvalues thatwereconsiderablymore negative han those isted in sev-eral recenttabulations.However, there is an error in theprocedure ollowed by Peryeaand Kittrick (1988) in the

calculation of the freeenergiesof the phases Hemingwayet al., 1989). The free energy of formation of Al3* re-ported by Hemingway and Robie (1977b, ncorrectly cit-ed previously as Hemingway et al., 1978) s basedon theassumption that gibbsite solubility was accuratelydeter-mined by Kittrick (1966). If the revised solubility forgibbsite is accepted, hen the free energy of formation ofAl3* must be recalculatedbecause he Gibbs free energyof formation of gibbsite has been determined by calori-metric methodsand represents he reference alue for Alin the calculation. The revised free energy of formationof Al3* would be -487 .5 kJ/mol and the correctedGibbsfree energiesof formation of corundum, boehmite, anddiasporewould be -1583.7, -919.1, and, 923.4 kJ /mol, respectively Hemingwayet al., 1989).These esults

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452are approximately -1.5 kJ/mol more negative than re-sults reported in recent tabulations (e.g., Robie et al.,r979).The free energyof formation of Al3* is subjectto ad-justment,as discussed bove, f acidpretreatments shownto result in the best solubility data for gibbsite in acidsolutions. Assuming the other solubility resultsgiven byPeryeaand Kittrick (1988) to also be the best solubilityvaluesfor corundum, diaspore,and boehmite in acid so-lutions, then the free energyof formation of Al3* can becalculatedfrom these data and the free energiesof for-mation of -1582.2 kJ/mol (CODATA, Cox, 1978),-922.9 kJ/mol and -918.4 kJlmol (Hemingwayet al.,1978), espectively, or corundum, diaspore,and boehm-ite. These alculations ield -486.8, -487.0, and -486.8kJlmol, respectively, or the free energyof formation ofAl3+. The valuesare essnegative han the valueobtainedfrom the calculationsbasedupon gibbsite solubility, butthe results agree within experimental error. If the solu-bility for gibbsitegiven by Bloom and Weaver (1982) isused n placeof the data from Peryeaand Kittrick (1988),one obtains -487.3 kJ/mol for the free energyof for-mation of Al3+. Whether or not the free energy or Al3*should be modified, these results demonstrate that thefree energies or the phases orundum, diaspore,boehm-ite, and gibbsite,as given above, are consistent.This, ofcourse,assumeshat eachphasehasequilibrated with thesame Al solution species.

Rnr,,lrrvr srABrlrry oF THE A(OH)3POLYMORPHS

Beforethe Gibbs freeenergyof formation of boehmitecan be calculated, he Gibbs free energyof formation ofAl3* or AI(OH);, or both, must be established.To dothis, the relativestability of the AI(OH), polymorphsmustbe established.If the evidencecited by Sanjuan and Michard (1987)is substantiated discussionn an earliersection), hen theconclusionsthey reached would directly follow and es-tablish the free energy of formation of AI(OH);. It isappropriate, herefore, o evaluate he supportingstudies.The results of Couturier et al. (1984) and Schoen andRoberson 1970)arecritical and are discussed elow.Theresult of Verdes and Gout (1987) supportseither view-point and thus is not definitive.Couturier et al. (1984) have reported that they mea-sured he stability constantsof hydroxocomplexesof Al3+and AI(OH); with oxalic acid. The authors chose thisprocedurebecause hey believed that only dissolved spe-cies would be involved in their study, thus eliminatingthe problem of identifying the controlling AI(OH), phasethat is necessaryn the application of solubility studies(e.g.,May et al., 1979;Hemingway, 1982).Using ther-modynamic properties or Al3* (Hemingway and Robie,1977b),Couturier et al. (1984)calculated he free energyof formation of AI(OH); as -l3ll.3 kJ/mol, a valuesubstantiallymore negative han the value of - 1305 kJ/

HEMINGWAY ET AL.: THERMODYNAMIC PROPERTIES OF BOEHMITEmol, calculatedby Hemingwayet al. (1978)and Hem-ingway 1982).Couturier et al. (1984) assumed hat the strong com-plexes hat occur between Al solution speciesand oxalicacid would prevent precipitation of AI(OH), phases.Cou-turier et al . (1984)alsoassumed hat no mixed hydroxyl-oxalate complexeswere formed. These assumptions ap-pear to be valid in acidic solutions(pH < 5), but may beinvalid in morebasicsolutions Sjdberg nd Ohman, 1985;Bilinski et al., 1986).Sjoberg nd Ohman (1985)studiedthe equilibria between Al solution species,hydroxide, andoxalic acid from pH 0.2 to approximately 7. The upperpH limit in their study (coincidently the pH region inwhich May et al., 1979,observedan offset to lower sol-ubility) was set by the onset of precipitation as deter-mined by turbidity measurements.Violante and Violante(1980) studied the effect ofpH and chelating organic an-ions on the synthesisof aluminum hydroxides and oxy-hydroxides.Oxalic acid was found to not inhibit bayeriteprecipitation in alkaline solutions at low concentrations,but as the concentration (with respect to Al) was in-creased,bayerite precipitation was inhibited and nords-trandite or gibbsite precipitated. The studies of Sjdbergand 6hman (1985)and Violanteand Violante(1980)showthat precipitation of A(OH)3 phasesdoes occur in thepresenceof oxalic acid in slightly basic solutions. Thus,the assumption made by Couturier et al. (1984) s invalidfor basic solutions and precipitation can be expected.Where precipitation does occur, the precipitation mech-anism will control the Al solution concentration and theequilibria with respect o oxalic acid will adjust accord-ingly. Therefore, the free energy of formation ofAI(OH); calculated by Couturier et al. (1984) must bequestioned.The second critical study cited by Sanjuan and Mi-chard (1987) was that of Schoenand Roberson(1970)who reported that they had observed a gradual disap-pearanceof gibbsite in solutions precipitating bayerite.Schoenand Roberson 1970)concluded hat bayeritewasmore stable than gibbsite in basic solutions. However,examination of the data presentedby Schoenand Rober-son (1970) suggestshat nordstrandite was misidentifiedas gibbsite.Schoenand Roberson(1970) identified earlyformed solids on the basisof one or two X-ray diffractionpeaks or calculatedd-values) hey considereddefinitive.Observedd-values were commonly from 4.7 to 4.9, andapproximately 4.4 and 2.2 A. Ttre d-values of 4.4 and2.2 wereassigned o bayerite and the 4.7-4.9 A d-valueswereassignedo gibbsite.Although bayerite exhibits a d-valueof 4.71 A, the early formed bayeritewasconsideredto have crystallizedwith poorly developed basalplanes.Although nordstrandite was not considered by Schoenand Roberson(1970), t appears o be likely becauseVio-lanteand Violante(1980)assigned -valuesof 4.72 Atobayerite, 4.79 A to nordstrandite, and 4.85 A to gibbsiteformed and examined under similar conditions.The discussiongiven above strongly questions he in-terpretations and results of Couturier et al. (1984) and

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HEMINGWAY ET AL.: THERMODYNAMIC PROPERTIES OF BOEHMITE 453Sanjuanand Michard (1987) with respect o the relativestabilitiesof the A(OH). polymorphsand their calculatedvalue for the free energyof formation of Al(OH);. How-ever, the interpretationsgiven abovesupport,but do notprove, the interpretations of Hemingway (1982), Hoveyet al. (1988), nd Apps et al. (1988).

Verdesand Gout (1988) provide evidence or the rel-ative stability of bayerite and gibbsite that is consistentwith the results of this study. On the basis of solubilitymeasurements, erdesand Gout (1988) conclude hatgibbsite is more stablethan bayerite and obtain a valuefor the Gibbs energyof formation of AI(OH); similar tothat proposedby Hemingwayet al . (1978)and Heming-way (1982). Further, using the free energyof AI(OH);,Verdesand Gout (1988)havecalculated 916 arrd-921kJlmol for the Gibbs freeenergyof formation of boehm-ite and diaspore, espectively, rom a combination of sol-ubility measurements and from crystallization fromamorphous oxides.The relative stability of the three common A(OH)3polymorphs also may be inferred from resultspresentedby Violante and Violante (1980) who studied the influ-enceofchelating organicanions on the synthesisofalu-minum hydroxidesand oxyhydroxides.The authors oundthat, in alkaline solutions, ncreasing he ratio of the com-plexing organicanion to dissolvedAl produceda changein the phase that precipitated, from bayerite to nord-strandite to gibbsite.This information is consistentwiththe inferencethat bayerite is the least stablepolymorphand gibbsite is the most stable.Sor,unrr,rry oF GTBBSTTEVarious studies have concluded that a range of freeenergieswill be shown by gibbsite samplesas a conse-quenceof differencesn crystallinity (e.g.,Helgesonet al.,1978;May et al., 1979;Bloom and Weaver,1982;San-juan and Michard, 1987). n acid solutions, here arethreefairly consistentdata sets hat may be representedasthreesubparallelcurves ofAl concentration vs. pH in the acidregion (seeFig. 5 of Hemingway, 1982).Frink and peech(1962)used the samegibbsitesample,but report solubil-ities that fall along the three curves (one ofwhich is de-fined by some of their data). Because he samegibbsite

samplewas used n theseexperiments, he crystallinity ofthe gibbsitecannot be the causeofthe observedsolubilitydifferences.Also, it is unlikely that equivalent degrees fcrystal imperfection could be obtained in differentgibbs-ite samples noteagreement f Kittrick, 1966 Singh, 1974;May et al., 1979).Therefore, it would appear that vari-ations in samplecrystallinity are not the major causeofobserveddifferences n gibbsitesolubility.It does not appearthat the acid pretreatment utilizedby Bloom and Weaver (1982) and Peryeaand Kittrick(1988) is the causeof the lower solubility observedbytheseauthors.Kittrick (1966),Singh 1974),and May etal. (1979, syntheticeibbsite)useddifferent gibbsite sam-plesbut obtained nearly identical solubility valuesfor Alin acid solutions. The solubility reported by Bloom and

Weaver(1982) for the treatedsample studied by Kittrick(1966) is lower, but it is in agreementwith results re-ported by Frink and Peech 1962)where he samegeneralexperimentalapproach wasused. Of greatestmportanceis the fact that Frink and Peech(1962), using the samegibbsite sample but different experimental parameters,found different solubilities, some of which agreed withthose of Kittrick (1966),Singh (197a),and May et al.(1979). Also, May et al. (1979) pretreatedtheir gibbsiresamplesby repeatedsuspension seven imes) in deion-ized HrO followed by centrifugation. Thus the sampleusedby May et al. (1979) is as likely to have had fine-grainedgibbsite particles removed and active surfacede-fects modified as that sample studied by Bloom andWeaver (1982).Several studies have shown that Cl- has an inhibitingeffect on the formation of crystalline AI(OH), (e.g.,Thomasand Whitehead, 93l; Hsu and Bates,1964;Hsu,1967;Turner and Ross, 1970;Rossand Turner, l97l).In the latter two studies, he concentrationsof mononu-clear and polynuclear Al ions were determined by theeight-quinolinolateextractionmethod (Turner, 1969)andthe amount of Al in the solid phasewascalculatedas hedifferencebetween he initial total of the dissolved Al andthe sum of the mononuclearand polynuclearspecies.Ofcritical importance was the observation that after ap-proximately 12d,, he concentrationof mononuclearspe-cies remainedessentiallyconstant or periodsof 100d ormore and the solid phase showed no evidence of anA(OH)3 phase.The solid phaseconsistedof a basic alu-minum hydroxychloride that was X-ray amorphous ex-cept at higher chloride ion concentrations and times ofapproximately 300 d. The period in which the concen-tration of mononuclear species remained constant in-creasedwith increasen the concentrationof Cl-. Follow-ing this period, gibbsiteappeared n the crystallinephase,the concentration of polynuclear speciesdecreasedo 0,and the concentration ofmononuclear species ncreased(e.9.,Fig. 28, Turner and Ross,1970).The differencesnthe solubilities ofgibbsite reportedby Bloom and Weaver(1982)and Peryeaand Kittrick (1988),as compared orhose of Kittrick (1966),Singh (1974),and May et al.(1979)are consistentwith the resultsofthese studiesandsuggest n alternative explanation o that ofacid pretreat-ment. The extendedperiod in which the concentrationofthe mononuclear species emains nearly constantand inwhich any aluminum hydroxychloride is X-ray amor-phous could easilybe mistakenly interpreted as showingequilibrium betweengibbsitesuspendedn such solutionsand the mononuclear pecies.The most probablecauseof the observeddifferences nAl solubility lies in differences n the experimental pro-ceduresused in the various studies. Recentadvances nsolution nuclear magnetic resonance NMR), in particu-lar the use ofthe Fourier-transform procedurebeginningin the 1970s,has allowed extensivedocumentation of thebehavior of Al in solutions, and that information is ap-plicable to this study. A major contributor in this area s

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45 4Akitt who, with hiscoworkers Akitt et al.,1972a,1972b,l98l ; Ak i t t andFarth ing,98la, l98lb, l98lc , l98ld,l98le; Akitt and Elders, 1985),has provided much ofthe information that will be drawn upon in the followingdiscussion.It has beenknown for many years hat Al-bearing so-lutions behavedifferently when the solutions have some-what differentchemistriese.g.,Hsu, 1967;RossandTur-ner,l97l1, Hemingway,1982and referencesherein;Tsaiand Hsu, 1984, 1985).Until recently, he natureof someofthese differenceshas been obscure.Akitt and Farthing(198 b) used solution NMR and gel-permeation hro-matography to study the Al speciespresent n two solu-tions preparedwith different procedures.Both solutionswereprepared o have a ratio (m) of OH/AI of 2.5. Bothsolutions were prepared at approximately 100 'C. Thefirst solution was preparedby hydrolysis of AlCl, by therapid addition of NarCOr. In the second,aluminum met-al was dissolved in an AlCl. solution. The first solutionshowed one peak in the NMR spectrum which was as-signed o the species AlO4Alrr(OH)ro(OHr),r]'*whichwill be describedby the usual symbol Alli. The secondsolution spectrum was more complex and interpretationby Akitt and Farthing (l98lb) suggested t least four Alspecies, wo of which were assigned o Al13* and[A(OH,)6]3* (designatedas Al3+). Akitt and Farthing(l98lb) indicated hat other hydrolysismethodsyieldedsolutions showing spectra that differed from those de-scribed above, but exhibited the same generalfeatures,that is, varying ratios of the sameAl species.Necessarily,these alternative hydrolysis methods involve changesnthe bulk chemistryof the solution [e.g., he useof AI(NO.),in placeof AlCl.l aswell as n the preparation procedures,but the work of Akitt and Farthing shows hat the differ-ent proceduresead to diferences n speciationof Al. Ak-it t et al . (1972b,p. 605) haveshown hat the concentra-tion of monomer Al speciess dependenton the procedurefollowed in the preparation of solutions with 0 < m

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(1962,1963)and of May etal. (1979) or solubilityof thenatural gibbsite sample and is consistentwith the pI!,Hemingway (I 982) postulatedfor nordstrandite.Using ,7Al solution NMR, the structure of Al in themore stable polymer proposed by severalauthors (e.g.,Tsai and Hsu, 1985)cannot be determined,nor can it beproved that a polymer rather than a crystallinematerialis present.However, work reportedby Botteroet al. (1980)has shown that AI(OH); and Allr* can be present n so-lutions preparedat 20 "C with m as low as 0.5. Conse-quently, additional studiesarenecessaryo determine hespeciesactually involved in the precipitation of AIOOHand AI(OH), phases.CoNcr,usroNs

The resultsdiscussedabove provide only a glimpseofthe complexity of the systemunder study. Chloride ionclearly interacts with Al and changes he mechanismofhydrolysis in a manner that is different from that of thenitrate ion. 27Alsolution NMR has not yet focusedontheseprocesses nd may not be able to resolve he struc-tural differences.However, the techniquehasestablishedbeyond doubt that the method of preparation of Al so-lutions may result in differences n the type and amountof speciespresentand, together with the preponderanceof other experimental data, supports the choice of thedata setsof Kinrick (1966),Singh 1974),and May et al.(1979) to determine the Gibbs free energies of Al3+(-489.8 + 4.0kJ/mol)and AI(OH); (- 1305.0+ t.3kJ/mol). Basedupon the measurements eported here, thisanalysis and the analysesof Hemingway et al. (1978),Hemingway 1982),Hovey et al. (1988),and Apps et al.(1988), the recommendedvalues for the entropy, Gibbsfree energy,and enthalpy of formation of boehmite are37.19+ 0.1J/(mol.K),-918.4 + 2.lkJ/mol,and, 996.4+ 2.2k1/mol, respectively.The recommendedGibbs freeenergyof formation of boehmite is intermediatebetweenthe value of -917.5 kJlmol calculatedby Apps et al.(1988) from solubility studies n basic solutions and thecorrectedvalue rom Peryeaand Kittrick (1988)of -919.1kJ/mol basedupon solubility studies n acidic solutions.The recommended esultsare consistentwith the anal-yses fHovey et al. (1988)and Apps er al. (1988); ow-ever, theseanalysesare not totally independent.Hoveyet al. (1988),using he Gibbs freeenergyof boehmite andset of solubilitydatarecommended y Apps et al. (1988,then in preparation),calculateda value for the Gibbs freeenergyof formation of AI(OH); of -1305.6 + 0.2kJ/mol. Apps et al. (1988)used he samesolubilitydata toestablish he Gibbs free energyof formation of boehmite(-917.5 kJlmol) and then calculateda consistent aluefor the Gibbs free energy of formation of Al(OH);(-1304.8 kJ/mol). Consequently, ifferencesn the in-terpretation of modelsemployedand model fit to exper-imental data result n small variations in the specificval-ues reported for either the Gibbs free energyof boehmiteor AI(OH);.Both differences n sample crystallinity and acid pre-

45 5treatment wereexamined aspossiblecausesor observeddifferencesn gibbsite solubility. Neither of these actorswere shown to be significant. Gibbsite would appear tocrystallize with an ordered and well-definedcrystal struc-ture for which a single value of the Gibbs free energy sappropriate. However, as with any mineral, grinding ofgibbsitemay result n distortion of the crystal surfaceandmay result in a surfaceenergycontribution in some ypesof measurements.Differencesobserved or the solubilityof gibbsite and commonly ascribed to differences ingibbsitecrystallinity or acid pretreatmentof gibbsitesam-ples more probably are causedby differences n the pro-ceduresused in the experiments which results n a finalstate or the Al that is different from that assumedby theinvestigator.

ACKNowLEDGMENTSWe wish to thank K. Wefers of Alcoa for preparing the two boehmitesamples and Helen V. Michel and Andrew W. Yee of Iawrence Berkeley

I-aboratory for the chemical analysis.We thank our U.S. GeologicalSur-vey colleagues Susan Russell-Robinson for help in the collection of DSCand X-ray data and John L. Haas, Jr. and J.J. Hemley for thoughtfulreviews of the manuscript. Constructive comments were prowidedby R.O.Sack,J.D. Hem, and J.Y. Walther, and they aregratefullyacknowledged.RnrnnnNcns crrED

Akitt, J.W., and Elders,J.M. (1985)Aluminum-27 nuclear magnetic es-onance studies of the hydrolysis of aluminum (III). Journal of theChemical Society,Faraday,81, 1923-1930.Akitt, J.W., and Farthing, AIan (l98la) Aluminum-27 nuclear magneticresonance studies of the hydrolysis of aluminum (III). Pan 2. Gel-prmeation chromatography. Joumal of the Chemical Society, Dalton,l 606- 1608.-(l98lb) Aluminum-27 nuclearmagnetic resonance tudies of thehydrolysis of aluminum (IID. Part 3. Stopped-flow kinetic studies.Joumal ofthe Chemical Society,Dalton, 1609-1614.-(l98lc) Aluminum-27 nuclear magnetic resonance tudiesof thehydrolysis ofaluminum (II!. Part 4. Hydrolysis using sodium carbon-ate.Joumal ofthe Chemical Society,Dalton, 1617-1623.-(l98ld) Aluminum-27 nuclear magnetic resonance tudiesof thehydrolysisofaluminum (IID. Part 5. Slow hydrolysis usingalurninummetal. Journal ofthe Chemical Society,Dalton, 1624-1628.-(l98le) Hydrolysis ofhexa-aquo-aluminum(III) in organicmedia.Journal ofthe Chemical Society,Dalton, 1233-1234.Akitt, J.W., Greenwood,N.N., and Khandelwal,B.L. (1972a)Alurninum-27 nuclear magnetic resonance studies of sulphato-complexes of thehexa-aquo aluminum ion. Journal of the Chemical Society, Dalton,1226-t229.Akitt, J.W., Greenwood,N.N., Khandelwal,8.L., and kster, G.D. (1972b)':7Al nuclear magnetic resonance studies of the hydrolysis and poly-merisation of the hexa-aquo-aluminum (III) cation. Journal of theChemical Society,Dalton, 604-6 10.Akitt, J.W., Farnsworth,J.A., and Letellier, Piene (1981)Nuclear mag-netic resonanceand molar-volume studies of the complex formed be-tween aluminum (III) and the sulphate anion. Journal ofthe ChemicalSociety,Faraday,81, 193-205.Apps,J.A., Neil, J.M., and Jun, C.-H. (1988)Thermochemicalpropertiesof gibbsite, bayerite, boehmite, diaspore, and the aluminate ion be-tween 0 and 350 "C. Iawrence Berkeley Laboratory-2l482, Berkeley,Califomia.Baes,C.F-, Jr., and Mesmer, R.E. (1981)The fiermodynamics of cationhydrolysis.American Journal of Science,281, 935-962.Barnhisel,R.I-, and Rich, C.L (1965)Gibbsite,bayerite,and nordstrand-

ite formation as affected by anions, pH and mineral surfaces.Soil Sci-enceSocietyof America Proceedings, 9, 53 1 534.Berman, R., Brown, T.H., and Greenwood, H.J. (1985) An internally

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HEMINGWAY ET AL.: THERMODYNAMIC PROPERTIES OF BOEHMITE