aluminum solubility control in different horizons of a podzol

DIVISION S-2—SOIL CHEMISTRY

Aluminum Solubility Control in Different Horizons of a PodzolM. Zysset,* P. Blaser, J. Luster, and A. U. Gehring

ABSTRACTAluminum extractability and solubility were investigated in detail

in six horizons of a Typic Haplohumod (FAO: Haplic Podzol) fromsouthern Switzerland. Pyrophosphate and oxalate extractions as wellas successive acid leaching indicated that in the Ah, (AE), and Bhhorizons reactive Al is mainly bound to soil organic matter, whereasin the Bs, BC1, and BC2 horizons it is of inorganic nature. In thelatter three horizons, infrared (IR) spectroscopy and transmissionelectron microscopy (TEM) revealed the presence of imogolite. Batchequilibrium experiments at 20°C in the pH range of approximately3.5 to 5.5 showed that the podzol profile can be divided into two partsof different Al solubility control. In the Ah and (AE) horizons, Alsolubility was found to be controlled by complexation reactions tosoil organic matter. Kinetic studies with samples of the Bh, Bs, BC1,and BC2 horizons showed that ion activity products with respect toboth Al(OH), and imogolite, (HO)3AI2O3SiOH, reached a constantvalue after reaction times of 16 d. For pH 4.1, the compilation ofall data revealed pAI + 0.5 pSi = 3.05 pH - 7.04 (r2 = 0.99) andpAI = 2.87 pH - 8.07 (r2 = 0.99). These data could be shown to beconsistent with either Al solubility control by imogolite-type material(ITM) with a log *KJ = 6.53 ± 0.09, which dissolves incongruently,or a simultaneous equilibrium with ITM and hydroxy-AI interlayersof clay minerals. For pH <4.1, data indicated solubility control by a1:1 aluminosilicate, e.g., poorly crystalline kaolinite.

IN THE LAST TWO DECADES, the anthropogenically in-duced acceleration of forest soil acidification has

been a topic of environmental concern. In many indus-trialized areas, the atmospheric deposition of sulfur andnitrogen compounds is a major source of proton inputto soils (see Matzner, 1992; Paces, 1985; van Breemenet al., 1984). In acidic soils at pH values <4.5, a majorpart of the acid input is buffered by dissolution of reac-tive Al phases. The released Al in soil solution can bephytotoxic and, thus, can pose a major ecological risk(see Sumner et al., 1991; Ulrich, 1989; Foy, 1984). De-spite this toxicity, Al solubility control in soils is notunderstood completely. In acidification models such asSAFE (Sverdrup et al., 1995), it is assumed that Alactivity in the soil solution is regulated by equilibriumwith an A1(OH)3 phase.

On the basis of both field measurements and labora-tory experiments several different solid phases havebeen proposed to control Al solubility in soils. Walker

M. Zysset, Dep. of Soil Sci., Swedish Univ. of Agric. Sci. (SLU),750 07 Uppsala, Sweden [present address: Jagerweg 6, 3014 Bern,Switzerland]; P. Blaser and J. Luster, Swiss Federal Institute for Forest,Snow and Landscape Research (WSL), 8903 Birmensdorf, Switzer-land; and A.U. Gehring, Institute of Terrestrial Ecology, ETH Zurich,Grabenstrasse 3, 8952 Schlieren, Switzerland. Received 23 Feb. 1998.*Corresponding author ([email protected]).

Published in Soil Sci. Soc. Am. J. 63:1106-1115 (1999).

et al. (1990) attributed Al solubility control in organicsoil horizons to exchange reactions between H and Alat soil organic matter. They found that equilibrium solu-bility was attained within hours and increased with de-creasing pH and increasing degree of saturation of soilorganic matter with Al. Their results agreed with earlierstudies by Cronan et al. (1986) and Bloom et al (1979).Mulder et al. (1989) concluded from results of a leachingexperiment that even in acid mineral soil horizons Alsolubility can be controlled by complexation with soilorganic matter. This finding was confirmed by soil solu-tion data obtained from acid brown forest soils in theNetherlands (Mulder and Stein, 1994). Berggren andMulder (1995) concluded in their study that, dependingon pH range, different phases can control Al solubility.They attributed Al solubility in acid mineral soil hori-zons from southern Sweden at pH <4.1 to complexationreactions with soil organic matter, and, at higher pHvalues, to an A1(OH)3 phase with a higher solubilitythan gibbsite. David and Driscoll (1984) found that soilsolutions from two Spodosol B horizons were often closeto equilibrium with synthetic gibbsite. In a laboratorystudy, Dahlgren et al. (1989) observed fast equilibriumof Spodosol Bs horizons with respect to A1(OH)3 fromboth conditions of undersaturation and oversaturation.They attributed control of Al solubility to hydroxy-AIinterlayers of expansible 2:1 layer silicates, since (i) noor only trace quantities of gibbsite could be detected inthe soil samples, and (ii) Al activities were oversaturatedwith respect to gibbsite. Similar results were reportedby Dahlgren and Walker (1993). Hydroxy-AI interlayersare less stable than gibbsite as was shown for smectites(Turner and Brydon, 1965; 1967).

Another phase that was proposed to control Al solu-bility in acid soils, is imogolite, a short range orderedaluminosilicate with an Al:Si ratio of 2:1. A less orderedmaterial with the same chemical composition is oftenreferred to as proto-imogolite allophane (Farmer et al.,1980), and the sum of imogolite and proto-imogoliteallophane as ITM (imogolite type material; Gustafssonet al., 1995). Imogolite was detected in podzols fromdifferent countries (see Tait et al., 1978; Farmer et al.,1980; Ross and Kodama, 1979; Childs et al., 1983; Gus-tafsson et al., 1995). However, not all podzols containimogolite (see Wang et al., 1986). According to Farmer(1987) the formation of imogolite and allophane re-

Abbreviations: AAS, atomic absorption spectroscopy; BS, base satu-ration; CEC, cation-exchange capacity; FIA, flow injection analysis;IAP, ion-activity products; ICP-AES, inductively coupled plasmaatomic emission spectroscopy; ICP-MS, inductively coupled plasmamass spectroscopy; IR, infared; ITM, imogolite-type material; TEM,transmission electron microscopy.

1106

ZYSSET ET AL.: ALUMINUM SOLUBILITY CONTROL IN A PODZOL 1107

quires a pH value >4.6 to 4.9. Farmer (1987) suggestedthat Al solubility in soils containing imogolite may becontrolled by a simultaneous equilibrium between anA1(OH)3 phase and imogolite. Dahlgren and Ugolini(1989) found evidence for this mechanism in soil solu-tions obtained from Bs and C horizons in a TephriticSpodosol. Data from batch equilibrium experiments us-ing spodic B horizons also suggested this mechanism(Su et al., 1995).

Most previous studies on Al solubility in soils investi-gated specific horizons, mainly A or spodic B horizons.Many dealt only with a narrow pH range, and this canlead to difficulties in assessing the solubility controllingphase. The aim of this research was to investigate theAl-solubility behavior in a complete sequence of hori-zons of an acidic soil profile in the approximate pHrange between 3.5 and 5.5, which is characteristic formany podzolized soils.

SAMPLES AND METHODS

Soil ProfileSamples were collected from a soil profile at Copera in

southern Switzerland (46°8' N, 8°59' E). The pedon was classi-fied as a Typic Haplohumod (Soil Survey Staff, 1996), HaplicPodzol (FAO, 1988), or as cryptopodzolic soil (Blaser et al.,1997). The soil developed on mica-rich gneissic bedrock underchestnut forest and is characterized by uniform incorporationof organic matter deep into the profile leading to only a weaklydeveloped, barely visible eluvial horizon.

Soil AnalysisField moist soil samples from each of the six horizons were

passed through a 2-mm sieve and stored at -20°C. Preliminarytests showed that freezing and melting did not affect the resultsof the experiments. All chemical analyses were performed inthree replicates. The mass of added soil samples to extractionsolutions were corrected for soil moisture content to obtainequal soihsolution ratios (mass:mass) for all horizons. The pHvalue was measured potentiometrically in deionized water(soil:solution ratio = 1:2; 30 minutes equilibration time). Ex-changeable cations were determined in 1 M NH4C1 extracts(soil:solution ratio = 1:10; 1-h extraction time) by inductivelycoupled plasma atomic emission spectrometry (ICP-AES; Op-tima 3000, Perkin Elmer Corp., Norwalk, CT). Exchangeableprotons were calculated by the difference of total and Al-induced acidity, which were determined according to Yuan(1959). Charge equivalents of exchangeable H, Al, Na, K,Ca, Mg, Mn, and Fe were used to estimate effective cation-exchange capacity (CECrft) and base saturation (BS). Reactiveforms of Al and Si were estimated by pyrophosphate (Alp)and oxalate extractions (Alox and Siox) using the methods ofBascomb (1968) and Schwertmann (1964), respectively. Theconcentrations of these elements in the extracts were deter-mined by flame atomic absorption spectrometry (AAS, PyeUnicam PU 9200, Unicam Ltd., Cambridge, UK).

A portion of the <2-mm fraction was dried at 60°C for48 h and used for the following analyses. After removingorganic matter by H2O2, particle size distribution was deter-mined by the pipet method described by Gee and Bauder(1986). Total Al concentrations (Al,01) in the samples weredetermined by analyzing Na-tetraborate digests dissolved indilute nitric acid with inductively coupled plasma mass spec-trometry (ICP-MS, VG PlasmaQuad PQ 2 Plus, VG Elemen-

tal, Winsford, UK). The carbon content was measured witha CN analyzer (Carlo Erba Instruments NA 1500 Series 2,Milan, Italy).

Leaching ExperimentTo test whether added hydrogen ions are mainly buffered

by the release of organically or inorganically bound Al, sam-ples were successively leached with a solution of the chemicalcomposition of 0.01 M HNO3 and 0.03 M NaNO3 and subse-quently extracted with oxalate or pyrophosphate as describedabove. Since it was intended to leach 40-50% of reactive Al,the number of teachings were varied for the different horizonsdepending on their contents of Alox: Ah: 5 leachings; (AE):6; Bh: 8; Bs: 6; BC1: 5; BC2: 4. The leachings were performedin duplicates in batch reactors with a soil: solution ratio of1:10. After a 15-h extraction on an end-over-end shaker at aconstant temperature of 20 ± 1°C, the suspension was centri-fuged for 15 min at 7000 g and the supernatant filtered (cellu-lose nitrate 0.45 jxm; Sartorius Corp., Edgewood, NY). Theremoved solution was replaced by fresh reaction solution andthe extraction procedure repeated. In all filtrates concentra-tions of Al were determined by AAS (Pye Unicam PU 9200).After the last leaching step, pyrophosphate or oxalate extract-able Al was determined in the solid residue as described aboveand referred to as Al* or AloX. For all horizons total leachedAl (A1L) differed less than 3% within two replicates.

Infrared Spectroscopy and TransmissionElectron Microscopy

Infrared Spectroscopy was performed on acid dispersibleclay fractions of Bs, BC1, and BC2 material. For the prepara-tion of the <0.5-|jun fraction, the method described by Farmeret al. (1980) was used. The IR spectra were obtained in trans-mission mode on pellets containing 1.5 mg of freeze-dried clayin 150 mg of a matrix of KBr. The pellets were dried at 150°Cin an oven for 12 h. Infrared spectra were recorded between4000 and 400 cm-' (Perkin Elmer 1760 X) and between 500and 200 cm"1 (Perkin Elmer 1700 X). In a second run, thesame pellets were heated at 400°C for 12 h and then anotherset of IR spectra was measured. The heating step was per-formed to distinguish between mineral phases in terms of theirthermal stability. For TEM, one drop of diluted suspensioncontaining the acid dispersable <0.5-u,m fraction was driedon a carbon-coated Cu grid. Analyses were performed witha JEM-2000 EX TEM (JEOL Ltd., Akishima, Japan).

Batch ExperimentsAl solubility behavior was investigated by batch experi-

ments on an end-over-end shaker at 20 ± 1°C. All reactionsolutions contained a constant ionic background of 0.03 MNaNO3 to reduce variation in ionic strength. To determine thetime required to reach equilibrium with inorganic Al phases,kinetic series were performed at selected HNO3 concentra-tions [(AE): 18.0 mM; Bh: 5.5 and 27.5 mM; Bs: 5.7 and16.7 mM; BC1: 2.9 and 9.7 mM; BC2: 2.9 and 9.8 mM) atsoiksolution ratios of 1:10 and reaction times varying between3 h and 35 d. At the end of the reaction time the suspensionswere centrifuged for 15 min at 7000 g. To avoid any impactof filters on H+-activity, pH values were measured in a smallportion of supernatant with a combination glass electrode.The remaining supernatant was filtered (Sartorius cellulosenitrate 0.45 (xm) and used for the following analyses. Labilealuminum (Allabilc), was operationally defined as Al quicklyreacting with 8-hydroxyquinoline, and was determined by themanual method described by Luster et al. (1993). Concentra-

1108 SOIL SCI. SOC. AM. J., VOL. 63, SEPTEMBER-OCTOBER 1999

tions of monomeric Si (Sim), were measured colorimetrically(Koch and Koch, 1964). Total concentrations of Al, Si, and Swere determined by ICP-AES (Perkin Elmer Optima 3000).

On the basis of the results of the kinetic series, a reactiontime of 30 d was chosen for the equilibration experiments.The reaction solutions contained HNO3 (0-55 mM) or NaOH(0-8.8 mM). In most experiments, a soil:solution ratio of 1:10was used. To evaluate a possible effect of this ratio, a fewadditional experiments were performed with a soihsolutionratio of 1:2.5. For all equilibrium experiments, concentrationsof Aliabji;, were determined by the flow injection analysis (FIA)method of Clarke et al. (1992) as modified by Berggren andSparen (1996) and concentrations of Sim by the FIA methodof Thomsen et al. (1983) as modified by Gustafsson et al.(1998). Results obtained with the corresponding manual andFIA methods differed less than 10% for Allabile and less than15% for Sim.

Chemical Speciation and Evaluation of AluminumSolubility Controlling Phases

To evaluate Al-solubility, free A13+ concentrations werecalculated from measured Aliab,k by the computer programALCHEMI 4.0 (Schecher and Driscoll, 1987). It was assumedthat AllaMe includes monomeric Al-hydroxo-complexes as wellas A1SO4

+ and AlH3SiOi+ (Berggren and Mulder, 1995). Al-organic and polynuclear Al-complexes are not included to anysignificant degree (Clarke et al., 1992). The influence of con-sidering A1SO4

+ on calculated Al3+-concentrations was testedin cases in which total S-concentrations were measured, assum-ing that all S was bound to SO4~. Free Al3+-concentrationsdiffered <1.3% when calculated with or without the consider-ation of SO4~. Therefore, A1SO+ was neglected in the specia-tion calculations.

Stability constants log PS, and log *Kj, as well as heats ofreaction AH? used in chemical equilibrium calculations arelisted in Table 1. Concentrations of A13+ were converted toactivities by the Davies equation. Since dissolved Sim occursas uncharged H4SiO4 in the investigated pH-range, its activitycoefficient was assumed to be 1. Logarithms of ion-activityproducts (log IAP) were calculated for A1(OH)3 and imogolite,(HO)3Al2O3SiOH, according to Eq. [la] and [Ib], where pH,pAl, and pSi stand for the negative logarithms of H+-, A13+-

Table 1. Stability constants and heats of reaction at 298 K and1 atm of species and solid phases used in equilibrium calcu-lations.

and H4SiO4-activities:

log IAP A1(OH)3 = 3pH - pAl [la]

Reaction

Dissolved species:AI3+ + H2O o AIOH2+ + H+

A13+ + 2 H2O *» A1(OH)2+

+ 2H +

AI3+ + 4 H2O <=> AI(OH)4+ 4 H"1"

Alw + H4SiO4 o AlH3SiO2+

+ H+

Solid phases:

Log p» AH? [kj/mol]

-5.00f 57.7f

-lO.lOf 122.5t

-22.70f 206.lt

-2.38J 66.6*log *K?

A1(OH)3 + 3 H+ « Al3

+ 3 H2O gibbsite 7.74§ -105.0§0.5 (HO)3Al2O3SiOH + 3H+

« Al" + 0.5 H4SiO4+ 1.5 H2O imogolite 6.00fl -96.8#

0.5 (HO)3AI2O3SiOH + 3H+

» AI3+ + 0.5 HjSiOi proto-imogolite+ 1.5 H2O____________sol________7.02# -96.8#

t Nordstrom and May (1996)t PokTovski et al. (1996)§ Palmer and Wesolowski (1992)H Farmer and Fraser (1982)# Liimsdon and Farmer (1995).

log IAP (HO)3 Al2O3SiOH = 3pH - 0.5 pSi - pAl

[Ib]Diagrams of pAl vs. pH and (pAl + 0.5 pSi) vs. pH were usedto evaluate possible phases which control Al solubility in theprofile. If Al activities are in equilibrium with A1(OH)3 orITM, the corresponding diagram reveals a linear relationshipwith the slope of 3 corresponding to the number of H+-ionsconsumed in the dissolution reaction, and the intercept to thenegative value of the stability constant, —log *KS. Further-more, at equilibrium log IAP = log *KS. In pH ranges in whichexperimental data suggested a linear relationship, the slopeand the intercept on the ordinate were estimated by structuralanalysis (see Webster, 1997), and 95% confidence intervalswere estimated according to Kendall and Stuart (1967). Bycontrast to the widely used linear regression, structural analy-sis takes uncertainties in both x and y values into account.From preliminary experiments, these were estimated to be±0.02 for pH, pAl, and pSi. To compare our results withliterature data, experimentally determined values of log IAPand log *KS were converted to standard conditions by thevan't Hoff equation and referred to as log IAP° and log *K°.

Neal et al. (1987) showed that, due to autocorrelation, dia-grams as used here can produce apparent order in data, evenwhen the data consists of random numbers, and this can leadto false interpretation of experimental results. Xu and Harsh(1995), however, concluded in their study that autocorrelationdoes not interfere with the interpretation of solid phase controlfor natural gibbsite at pH <5.5.

In addition to Al solubility control by A1(OH)3 or ITM,two special cases were tested: (i) simultaneous equilibriumwith ITM and an Al-hydroxide, and, (ii) congruent dissolutionof ITM. The dissolution of an Al-hydroxide can be describedby Eq. [2a] and, in equilibrium, pAl vs. pH is given by Eq.[2b], where log *KA is the stability constant.

A1(OH)P-Z)+ + zH+ «* A13+ + z H20 [2a]pAl = z • pH - log *KA [2b]

The dissolution of an aluminosilicate can generally be de-scribed by Eq. [3a], and pAl vs. pH in equilibrium by Eq. [3b],with log *K[ being the stability constant.

o x A13

y H4SiO43x + w — 4y H20

pAl = 3 • pH - y- • pSi - - • log *K,x x

Pa]

[3b]

If Al is in equilibrium with both phases, the right hand sidesof Eq. [2b] and [3b] become equal, and, therefore, pSi can becalculated according to Eq. [4].

pSi = -(3 - z) • pH + - • log *KA - - • log *K, [4]y y yFor the derivation of the pAl vs. pH relationship assuming

congruent dissolution of ITM, stoichiometric stability con-stants C*K° (solubility of ITM, see discussion section), CK] (for-mation of AlH3SiO^; Table 1), and cpn (mononuclear 1:1,1:2,and 1:4 Al-hydroxo-complexes; Table 1) were used. Equations[5a] and [5b] must be fullfilled, where [] denote concentrationsand [Al],otal and [Si]lotai stand for total concentrations of dis-


Table 2. Chemical parameters of the fraction <2 mm in the six horizons of the profile; pH value measured in H2O; content oforganic C (COIg); exchangeable cations; effective cation exchange capacity (CECe(t); base saturation (BS); standard deviation (SD) ofthree replicates. ________ ________

Horizon Sample pH SD Al Ca Fe Mg Mn Na H CECeB SD BS SD

Ah(AE)BhBsBC1BC2

cm3-78-14

18-2850-6090-100

120-135

H2O4.274.674.794.895.155.49

190.497.863.113.94.12.8

£

1.20.41.00,20.00.0

114.669.131.212.78.73.4

13.41.50.60.20.20.3

0.20.00.00.00.10.2

2.81.00.10.20.20.2

3.51.00.30.10.20.2

iol</kg —0.20.00.00.00.10.0

0.10.60.30.10.20.2

10.64.72.4<1<1<1

145.477.934.913.49.54.4

4.96.71.30.40.60.1

a,'l

13.65.23.74.87.1

19.2

0.20.00.20.21.12.0

solved Al and Si, respectively.

<*K? = 1̂ 1[H+]3

[Si]total= 0.5 • [Al]total

[5a]

[5b]It is assumed that [Al],olai consists of the species [A13+], mono-meric 1:1, 1:2, and 1:4 Al-hydroxo-complexes and[AlH3SiO5+], and [Si]to,a, is the sum of [H4SiO4] and

- Thus, [Al]tocal and [Si]totai can be calculated ac-cording to Eq. [6a] and [6bj.

[6a]

[6b]

Equations [6a] and [6b] are put into Eq. [5b], which is solvedfor [H4SiO4] to replace this species in Eq. [5a]. After rearrang-ing, Eq. [7] is obtained.

0.5 •

Lrtljtotai —

rsii. . -

|rtl J 1 ^ LJH J

n = l,2,4 [ii \

CK! • (C*K?)2 • [H+]5

[AP+]

n-r.sio.i 4- ' ' (c s-* ' 1- J

J

C*K!> • [H + ]3

1= 0 [7]

For a certain H+-concentration, Eq. [7] can be solved for A13+

by using, e.g., the Newton-Raphson method. This was donefor several pH values between 3.5 and 5.0, assuming an ionicstrength of / = 0.04, which is similar to the mean / determinedin the equilibrium experiments, and a temperature of 20°C.The corresponding plot pAl vs. pH revealed a slope of 2.0.

RESULTSPhysical and Chemical Properties

of the Soil SamplesThe particle size distribution was dominated by the

sand fraction with relative contents between 51 and 62%

in the different horizons. The contribution of the siltfraction was in the range 28 to 37%. The clay contentshowed a decrease with depth from 17 to 6% (datanot shown). The pH determined in H2O increased withdepth from 4.27 to 5.49 (Table 2). Organic C exhibiteda decrease within the profile by almost two orders ofmagnitude. The effective cation exchange capacity de-creased from 145 mmolc/kg soil in the Ah horizon to4.4 mmolc/kg in the BC2 horizon. The base saturationvaried between 3.7 and 19.2% with a minimum in theBh horizon. Selective extractions showed for both Alpand Alox a maximum in the Bh horizon (Table 3). In thehorizons in which organic matter represented a majorcomponent, Alp was similar to Alox. In the lower hori-zons, the Alp decrease was more pronounced than theone of Alox. The ratio Alox/Altot varied between 0.06 and0.16. The (Alox - Alp)/Siox ratio as a measure to estimatethe Al/Si ratio in amorphous Al silicates (see Parfittand Kimble, 1989) varied between 0.68 and 3.32 withinthe four lower horizons. The low value observed for theBh horizon is uncertain as indicated by a high standarddeviation. The contribution of exchangeable Al, (Alexc),to Altot was =1% in the Ah and even smaller in theother horizons.

Leaching ExperimentThe relative amounts of A1L varied for the different

horizons between 39 and 58% of Alox (Fig. la). Oxalateextractable Al and (A1L + AloX) showed the same distri-bution pattern within the profile. However, the latterwas slightly higher in all horizons. A different behaviorwas observed for Alp. In the three uppermost horizons,(A1L + Alp) and Alp revealed the same distributionpattern with similar values (Fig. Ib). In the three lower-most horizons, however, Alp and Alp were similar and(A1L + Alp1) were significantly higher than Alp.

Spectroscopic and Microscopic AnalysesIn the range 3500 to 3800 cm'1 the IR spectra of the

acid dispersable clay fraction of the Bs, BC1, and BC2Table 3. Contents and ratios of different Al and Si fractions in the soil profile; p: pyrophosphate extractable; ox: oxalate extractable;

exc: exchangeable; tot: total; SD: standard deviation of three replicates; nd: not determined.Horizon

Ah(AE)BhBsBC1BC2

Alp

2523564871314633

SD Alox

7 2527 376

31 4993 3811 1951 191

SD

261

1225

Alp/Alm SD Sira SD (Alra - Ay/Si. SD

mmol/kg1.000.950.980.340.240.18

0.030.020.060.010.010.01

<10<10

18754860

__1222

__

0.683.323.122.64

1.760.190.120.12

Ale«

38.223.010.44.22.91.1

SD Al,,,,

1.3 24662.1 28850.4 30840.1 33090.1 33850.1 nd


100Al [mmol/kg]

200 300 400 500 600

20 -

40 -

60 -

80 -

100 -

120 -

140 -

100Al [mmol/kg]

200 300 400 500 600

20- b)

40 •

£ 60 -

80 -

100 -

120 -

140 -Fig. 1. Results of acid leaching experiment for samples from different

depths in the profile (Table 2); a) oxalate extraction; b) pyrophos-phate extraction; A I,,, and Alp:oxalate and pyrophosphate extract-able Al in untreated samples; A1JX and Alpioxalate and pyrophos-phate extractable Al in solid residues after successive acid leaching;AlL:total Al removed by acid leaching; A1L + Al?x and A1L +AIR:siim of total leached Al and extractable Al in solid residues.

horizons dried at 150°C exhibited bands at 3698 and 3627cm"1 (data not shown). These bands were attributed tostretching vibrations of structural OH groups in kaolin-ite (Farmer, 1974). Between 600 and 300 cm"1 fourbands at 536, 471, 428, and 343 cm"1 were found forsamples dried at 150°C (Fig. 2). Considering the bandsat 3698 and 3627 cm^1, which were typical for kaolinite,the bands at 536 and 471 cm"1 could be assigned unam-biguously to Si-O deformations and octahedral sheetvibrations of this mineral. Upon heating the samples to400°C, the bands in the near and mid infrared range didnot change significantly. The bands at 428 and 343 cm"1

BC2 400 °CBC2150°CBC1 400 °CBC1 150°CBs 400 °CBs 150°C

200

w600 400

crrr1

Fig. 2. IR-spectra in the range 600 to 300 cm"1 of the acid dispersibleclay fraction • 0.5 (xin from the horizons Bs, BC1, and BC2 heatedat 150°C and at 400°C, respectively.

Fig. 3. Transmission Electron Microscopy of the <0.5-(xm acid dis-persible clay fraction of the BC1 horizon.

were characteristic for kaolinite as well as for imogoliteor proto-imogolite allophane (see Farmer et al., 1979;Russel and Fraser, 1994; Gustafsson et al., 1995). Afterheating at 400°C, the band at 428 cm"1 changed to aless pronounced shoulder and the one at 343 cm"1 de-creased in intensity.

In the samples from the Bs, BC1, and BC2 horizons,TEM revealed fibers with lengths of several 100 nmand diameters of approximately 10 nm (Fig. 3). Thismorphology was attributed to imogolite (see Farmer etal., 1980; Dahlgren, 1994).

Batch ExperimentsIn all kinetic series, strong initial increases in Aliabjie,

pH, and Sim were observed, which lasted up to 2 d(Al,abile), 5 d (pH), and 10 d (Sim) (data not shown).Subsequently, only little changes occurred. After 16 dof reaction time, log IAP with respect to both A1(OH)3(Eq. [la]) and imogolite (Eq. [lb]) remained practicallyconstant in all series with the exception of the (AE)horizon. Selected examples are shown in Fig. 4a and 4b.The same qualitative behavior was found for all otherkinetic series.

With respect to the solubility behavior of Al deter-mined in batch equilibrium experiments, the profilecould be subdivided into two parts. In the Ah and (AE)


10.0 • a)

0.0

8.0 -i

0.015 20time [days]

x(AE) ABh BBs »BC1 * BC2Fig. 4. Values of log IAP vs. time for selected kinetic series, a) log

IAP for Al(OH), (Eq. [la]); b) log IAP for imogolite (Eq. [lb]).Initial H+-concentrations were 18.0 mM for (AE), 27.5 mM forBh, 5.7 mM for Bs, 9.7 mM for BC1, and 9.8 mM for BC2.

horizon, pAl vs. pH exhibited a behavior which differedstrongly from the one equilibrated with an A1(OH)3phase (Fig. 5a). In the Ah horizon filtrates were under-saturated with respect to gibbsite at low pH values,whereas an oversaturation was observed for filtrateswith pH >4.9. A qualitatively similar behavior was ob-served for the (AE) horizon, where filtrates were under-saturated with respect to gibbsite at pH <4.3 and over-saturated at higher pH values. For 3.59 < pH < 4.85,a linear relationship of pAl vs. pH with a slope of 2.16was observed.

The four lower horizons differed distinctly in Al-solu-bility behavior from the two uppermost horizons (Fig.5a). The diagram of pAl vs. pH suggested a linear rela-tionship for pH >4.0, where all filtrates were oversatur-ated with respect to gibbsite. At lower pH values, pAlvs. pH curves leveled off and at pH <3.7 filtrates wereundersaturated with respect to gibbsite. Within the Bsand BC1 horizons, pAl vs. pH behaved equally for bothsoil:solution ratios of 1:10 and 1:2.5 at pH >4, whereaspAl was slightly higher in the 1:10 suspensions at lowerpH values. To allow a direct comparison with the otherion activity vs. pH plots presented later [pSi vs. pH; (pAl+ 0.5 pSi) vs. pH], structural analysis was performedfor pH >4.1, where all three plots exhibited a linearrelationship. Slopes of pAl vs. pH varied between 2.72and 2.99 (r2 > 0.99) for the different horizons (Table4), and the compilation of all data yielded Eq. [8].

pAl = 2.87 • pH - 8.07 r2 = 0.99 [8]Lower and upper confidence limits of the slope at the95% level were 2.75 and 2.99, respectively. Mean values

7.0 -

6.0-

5.0 -

' 4.0 -

3.0 -

2.0

a) Gibbsite

Imogolite

3.0 3.5 4.0 4.5

pH

5.0 5.5 6.0

4.5 -

4.0 -

w 3.5 -Q.

3.0 -

2.5

Quartz

3.0 3.5

8.0 -

in 6.0 -ojt-

Q. 4.0 -

C)

4.0

Imogolite

4.5pH

5.0 5.5 6.0

Protoimogolite Sol

3.0 3.5 4.0 4.5PH

5.0 5.5 6.0

x Ah 1:10 x(AE)1:10 *Bh1:10 «Bs1:10oBs 1:2.5 • BC11:10 «BC21:10 o BC2 1:2.5

Fig. 5. Results of batch equilibrium experiment, a) pAl vs. pH; theimogolite line assuming congruent dissolution obeys Eq. [7]; b)pSi vs. pH; c) (pAl + 0.5 pSi) vs. pH. Closed symbols, soihsolution =1:10; open symbols, soilisolution = 1:2.5. The linear fittings in a)and c) were performed by structural analyses (see Webster, 1997)including all data at pH 4.1 of the Bh, Bs, BC1, and BC2 horizons.Structural analyses in b) was performed separately for both pH• 4 . 1 and pH 4.1 for each of the four lower horizons. Solubilitiesof gibbsite (Palmer and Wesolowski, 1992), imogolite (Farmer andEraser, 1982), proto-imogolite sol (Lumsdon and Farmer, 1995),quartz, and amorphous Si(); (Stumm and Morgan, 1981) werecalculated for 20°C by the van't Hoff equation.

of log LAP0 A1(OH)3 (Eq. [la]) varied between 8.25 and8.37 for the different horizons and the compilation ofall data yielded log LAP0 = 8.35 ± 0.09 (Table 4).

In many filtrates dissolved Sim was oversaturated withrespect to quartz, and in two filtrates at low pH valuesa slight oversaturation with respect to amorphous SiO2was observed (Fig. 5b; thermodynamic data fromStumm and Morgan, 1981). In the Ah and (AE) hori-zons, a slight increase of pSi vs. pH was found over the


investigated pH range. In the four lower horizons, pSiincreased strongly up to pH = 4.1 with slopes varyingbetween 1.23 and 1.76. At higher pH markedly lowerslopes between 0.25 and 0.40 were observed. In contrastto this qualitative similarity, the pSi vs. pH curves exhib-ited quantitative differences with respect to pSi decreas-ing in the order Bh > Bs > BC1 > BC2. The soil:solution ratio had only little effect on pSi.

The (pAl + 0.5 pSi) vs. pH data for the Ah and (AE)horizons differed markedly from the behavior of ITM(Fig. 5c). By contrast, (pAl + 0.5 pSi) for the four lowerhorizons increased with pH exhibiting slopes between2.87 and 3.17 at pH >4.1 (r2 > 0.99) (Table 4). Thecompilation of all data yielded Eq. [9].pAl + 0.5 • pSi = 3.05 • pH - 7.04 r2 = 0.99 [9]Lower and upper confidence limits of the slope at the95% level were 2.91 and 3.20, respectively. Mean valuesof log LAP0 (imogolite) (Eq. [Ib]) varied between 6.43and 6.66 for the different horizons and the compilationof all data yielded log IAP° = 6.53 ± 0.09 (Table 4).

DISCUSSIONFirst the extractability of Al in the different horizons

of the Typic Haplohumod is discussed. Since Alexc ismuch lower than Alox or Alp, the reactive part of totalAl is mainly associated with amorphous and/or poorlycrystalline inorganic phases or bound to organic matter.Pyrophosphate is known to extract preferentially Althat is coordinated with soil organic matter (see Parfittand Childs, 1988), whereas oxalate dissolves in additionAl in amorphous and poorly crystalline inorganic phasesas well as Al from chloritized vermiculite (Fordham andNorrish, 1983). On the basis of this, the podzol profilecan be subdivided into two parts of different extractionbehavior. In the three upper horizons with high contentsof Corg, Alp = Alox indicates that most of the reactive Alis associated with organic matter. In the three lowerhorizons which have low contents of Corg, the A1P/A1OXratios in the range 0.18 to 0.34 provide evidence thatthe major part of reactive Al is bound to amorphous orpoorly crystalline inorganic compounds.

In the four lower horizons, Siox is indicative of theoccurrence of allophane or ITM (see Farmer et al., 1983;Dahlgren, 1994; Gustafsson et al. 1995). Furthermore,ratios of (Alox — Alp)/Siox between 2.6 and 3.3 in the Bs,BC1, and BC2 horizons indicate that these phases arerich in Al.

A comparison of (A1L + Al£x) with Alox in the leaching

experiment shows that added acid mainly dissolves oxa-late extractable Al. The slightly higher values of (A1L +AlSj) compared with Alox suggest that minor amountsof crystalline Al bound to primary or secondary mineralsare also dissolved. These findings are in good agreementwith a study by Dahlgren and Walker (1993), whoshowed in kinetic experiments using spodic Bs horizonsthat acid mainly dissolves oxalate extractable Al. In theAh, (AE), and Bh horizons similar values of (A1L +Alp) compared with Alp as well as with Alox provideevidence that added acid mainly dissolves Al bound toorganic matter. In the three lowermost horizons, Alp* issimilar to Alp indicating that Al associated with inor-ganic amorphous or poorly crystalline compounds is themajor source of acid leachable Al. These findings areconsistent with an investigation by LaZerte and Findeis(1995) who stated that Alox is most affected by acidicleaching when the ratio (Alox - Alp)/Alp is above 0.3 to0.7 and Alp becomes most important at lower ratios.

The IR spectra reveal kaolinite in Bs, BC1, and BC2horizons. The occurrence of this mineral is proven bythe absorption bands at 3698, 3627, 536, and 471 cm"1.Imogolite or proto-imogolite allophane have character-istic bands at 428 and 348 cm"1 where they overlap withbands of kaolinite. The absorption bands which canunambiguously be assigned to kaolinite show no signifi-cant change in intensity after heating to 400°C. By con-trast to kaolinite, imogolite as well as proto-imogoliteallophane are unstable at this temperature. Thus, thedecrease in intensity of the two bands at 428 and 348cm"1 is indicative of these phases. The detection of athread-like morphology by TEM analysis provides fur-ther evidence for the occurrence of imogolite in the Bs,BC1, and BC2 horizons.

With respect to Al solubility as determined in batchexperiments, the profile can be subdivided into twoparts consisting of the (i) Ah and (AE) and (ii) Bh, Bs,BC1, and BC2 horizons. In neither the Ah nor the (AE)horizons, was equilibrium attained with respect toA1(OH)3 or ITM within 30 d. The observed behaviorrather suggests that Al solubility is controlled by com-plexation reactions to soil organic matter. In such a case,Al solubility is dependent on the amount of reactive Alin the system (see Walker et al., 1990; Berggren andMulder, 1995). Cronan et al. (1986) and Walker et al.(1990) found in batch experiments with organic horizonsthat solutions were undersaturated with respect togibbsite at pH values lower than 4.4 to 4.8, the exactvalue depending on the amount of reactive Al, whereas

Table 4. Results of batch equilibration experiments for the horizons Bh, Bs, BC1, and BC2 as well as the compilation of all data: pH-range; number of data points n; slope of pAl vs. pH (Al-hydroxide) or (pAl + 0.5 pSi) vs. pH (ITM), respectively, obtained fromstructural analysis; correlation coefficient r2; values of log IAP" were calculated by Eq. [la] or [Ib] and subsequently converted tostandard conditions by the van't Hoff equation; standard deviation, SD.

Horizon

BhBsBC1BC2All

PH

4.16-4.954.18-5.024.16-4.884.21-4.884.15-5.02

n

61245

27

slope r2 log IAP° SD slope r2

A 1 U J -J

2.892.912.992.722.87

1.001.001.001.000.99

8.378.408.258.308.35

0.050.060.030.130.09

3.023.093.172.873.05

1.001.001.000.990.99

log IAP°

6.436.556.496.666.53

SD

0.040.050.060.100.09


an oversaturation was observed at higher pH. Our re-sults from the Ah and (AE) horizons are consistent withthese findings, and, therefore, it is concluded that Alsolubility is controlled by complexation to soil organicmatter. This mechanism is consistent with the results ofthe selective extractions and acid leaching experimentswith Ah and (AE) material.

In the Bh, Bs, BC1, and BC2 horizons dissolved Alattained equilibrium after 16 d in all kinetic series basedon stable values of log TAP for A1(OH)3 and imogolite,suggesting that Al solubility is controlled by an inorganicphase. The simultaneous rapid initial increases in pH,Allabiie, and Sim in all kinetic series provide evidence thatadded acid is mainly buffered by release of Al and thata reactive aluminosilicate is involved in proton buff-ering processes.

As shown in Fig. 5a, the Bh, Bs, BC1, and BC2 hori-zons exhibit very similar pAl-pH reationships which sug-gest Al solubility control by the same solid phase. Con-sidering pAl vs. pH, the mean log IAP° of 8.35 forA1(OH)3 (Eq. [la]) at pH >4.1 is similar to values re-ported in the literature. Gustafsson et al. (1998) foundlog LAP0 = 8.29 for spodic B horizons containing imogo-lite from central Sweden and Finland. Dahlgren et al.(1989) reported a lower value of 8.1 for Spodosol Bshorizons in the northeastern USA containing no or onlysmall amounts of imogolite. Su et al. (1995) determinedlog IAP° = 8.73 ± 0.16 for Bs and BC horizons of aTephritic Spodosol which contained imogolite as a ma-jor component in the clay fraction. In this study, statisti-cal analysis revealed a slope of pAl vs. pH <3 at the95% confidence level, suggesting that Al solubility maynot be adequately described by assuming equilibriumwith A1(OH)3. Dahlgren and Walker (1993) found inbatch experiments with Spodosol Bs horizons a slopeof 2.7 between pH 3.1 and 5.2. On the basis of a studyby Bloom et al. (1977), these authors attributed Al-solubility control to hydroxy-Al interlayered 2:1 layersilicates arguing that if the hydroxy-Al interlayer phaseexhibits a positive charge to compensate the permanentnegative charge of the clay mineral, a slope of less than3 is obtained. In our study, Al-solubility control by chlo-rite or hydroxy-interlayered vermiculite, which are ma-jor mineral compounds in the investigated soil profile(Eggenberger, 1995), could lead to a slope of less than3 in the pAl vs. pH diagram.

Structural analysis of (pAl + 0.5 pSi) vs. pH for thefour lower horizons at pH >4.1 yields a slope of 3.05which, at the 95% level, does not significantly differfrom the theoretical value of 3.00 for ITM. Therefore,the mean log IAP° of 6.53 ± 0.09 can be considered torepresent log *K!? of ITM in the four lower horizons.This value is within the range of reported stability con-stants for synthetic imogolite (log *K° = 6.00, Farmerand Fraser, 1982; log *K» = 6.52, Su and Harsh, 1994)and for proto-imogolite sol (log *K° = 7.02, Lumsdonand Farmer, 1995). Thus, the consideration of (pAl +0.5 pSi) vs. pH suggests Al-solubility control by ITM.The uniform increase of pSi from pH = 4.1 up to pHvalues significantly higher than 5 may indicate that Al-solubility control by ITM is also valid in the pH range

5 to 6, where Aliawie was below detection limit. Thereare few studies on selected podzol horizons in whichindications for Al-solubility control by ITM were found.For a spodic Bs horizon, Gustafsson et al. (1998) re-ported log *K° = 6.64, and for Bs and BC horizons ofa Tephritic Spodosol log *K°S = 6.88 was found (Su etal., 1995). For the Bs, BC1, and BC2 horizons in thisstudy, Al-solubility control by ITM is consistent with theresults of selective extractions, leaching experiments, IRspectroscopy, and TEM. While Siox provides evidencefor the presence of low amounts of ITM also in theBh horizon, selective extraction and acid leaching dataindicate that reactive Al in this horizon is mainly boundto Al-organic phases. Thus, at pH <4.16, where under-saturation with respect to ITM occurs, Al solubility maybe controlled by complexation to soil organic matter.

As shown above, the independent considerations of(pAl + 0.5 pSi) vs. pH and pAl vs. pH lead to differentimplications on the Al solubility control in the fourlower horizons. However, there are two possibilities fora combined interpretation. As a first possibility, bothrelationships can be explained by solubility control byITM, which dissolves incongruently. Solubility controlby congruent dissolution of ITM is very unlikely, sincethis would lead to a pAl vs. pH relationship with a slopeof 2.0 as shown in Fig. 5a. Additional Al release by Al-organic phases might increase this slope. Consideringthe contribution of Alp to A1L in the acid leaching experi-ments, a significant effect of such a mechanism cannotbe excluded for the Bh horizon. In the Bs, BC1, andBC2 horizons, however, reactive Al bound to organicmatter very likely has only a minor effect on pAl vs. pH.The assumption of solubility control by incongruentlydissolving ITM is supported by data from the followingstudies. Su and Harsh (1994) reported that solutions inequilibrium with a synthetic imogolite were oversatur-ated with respect to gibbsite. Lumsdon and Farmer(1995) investigated the solubility characteristics of a syn-thetic proto-imogolite sol. Stuctural analysis of theirdata (runs D1-D4 and F1-F5) reveals slopes of 2.92,2.36, and 1.14 for (pAl + 0.5 pSi) vs. pH, pAl vs. pH, andpSi vs. pH, respectively. The findings of these studiessuggest that in presence of a short-range ordered alumi-nosilicate (i) Al activities can be oversaturated withrespect to gibbsite and (ii) a slope of pAl vs. pH between2 and 3 can occur.

As a second possibility, A13+ is assumed to be in asimultaneous equilibrium with ITM and hydroxy-Al in-terlayer. The latter phase has to be of the formAl(OH)^n)+ with n = 0.13 as indicated by the slope ofpAl vs. pH. According to Eq. [4], this would lead to aslope of pSi vs. pH of 0.26, which is within the observedrange of 0.25 to 0.40. Thus, our data are consistent withrespect to this mechanism which has been proposedearlier in the literature. Farmer (1987) proposed anequilibrium of this type for podzol Bs horizons. Dahl-gren and Ugolini (1989) found evidence for such anequilibrium in soil solutions from a Tephritic Spodosol,and this finding was confirmed in a laboratory studyconducted by Su et al. (1995).

Finally, the pH range below 4.1, where a break in pSi


vs. pH occurs, is discussed for the three lower horizons.In all these horizons, the equilibrium solutions weredistinctly undersaturated with respect to ITM atpH <3.8. This behavior can be explained by either kinet-ically impeded dissolution of Al, or control of Al activi-ties by a solid phase which is more stable than ITM. Inthe kinetic series with Bs material and a final pH valueof 3.70, log IAP with respect to both A1(OH)3 and imo-golite did not significantly change after 16 d, indicatingthat Al activities were in equilibrium with an inorganicsolid phase. Thus, it is likely that solutions attainedequilibrium at pH <4.1. The significant higher slopesof pSi vs. pH in this pH range when compared with thepH range above 4.1 may indicate Al solubility controlby another aluminosilicate than ITM. To evaluate thispossibility it was assumed that filtrates in the pH range3.58 to 4.07 (n = 14) were in equilibrium with an alumi-nosilicate of the general formula AlSiy(OH)wO((3+4y.w)/2),the solubility constant log *KS of which can be calculatedaccording to Eq. [10].3 • pH - pAl - y • pSi = log *KS = constant [10]

The value of y was varied until the slope -of log *KSvs. pH was close to zero. A value of 1.05 was determinedfor y and log *KS was calculated to be 5.04 ± 0.14. Theseresults suggest that an aluminosilicate with an Al:Si ratioclose to 1:1 controls Al solubility at pH values <4.1. Apossible candidate is kaolinite, which was detected byIR spectroscopy. Kittrick (1970) pointed out, that thestability of kaolinite in soils might vary between thecrystalline form and the one of halloysite. Thus, log *KSfor 0.5 Al2Si205(OH)4 + 3 H+ = A13+ + 2 H4SiO4 + H2Omight vary between 3.1 and 5.4 at 20°C (thermodynamicdata reported by Manley et al., 1987; Stumm and Mor-gan, 1981; Kittrick, 1969). The value estimated here iswithin this range, and, therefore, it is possible that Alsolubility at pH <4.1 is controlled by poorly crystallinekaolinite. It must be stated, however, that no directevidence for its presence could be found.

CONCLUSIONSThe first study on the Al solubility control for a com-

plete set of horizons of a podzol over the entire ecologi-cally significant pH range led to the following con-clusions:

1. Specific extraction and acid leaching experimentssuggest that reactive Al is mainly bound to organicphases in the Ah, (AE), and Bh horizons and toinorganic solid phases in the Bs, BC1, and BC2horizons. In the latter three horizons, oxalate ex-traction, IR, and TEM data indicate ITM as reac-tive inorganic Al phase.

2. Batch equilibrium experiments provide evidencethat in the Ah and (AE) horizons Al solubility iscontrolled by complexation reactions to soil or-ganic matter over the entire investigated pH range.This behavior is in agreement with specific extrac-tion and leaching data.

In the Bh, Bs, BC1, and BC2 horizons Al solubil-ity at pH > 4.1 can be explained either by an

equilibrium with incongruently dissolving ITM orby a simultaneous equilibrium with both ITM andhydroxy-Al interlayers of clay minerals. The databelow pH 4.1 for the three lower horizons suggestsolubility control by a 1:1 aluminosilicate, e.g.,poorly crystalline kaolinite.

ACKNOWLEDGMENTSThe authors thank the following persons at WSL: A. Diarra,

U. Beutler, R. Luescher, Dr. S. Zimmermann and K. Siegristfor technical support in the laboratory and at the field site,V. Michellod for performing selective extractions as well asleaching and kinetic experiments, and the staff of the centrallaboratory for performing ICP-AES analyses. Dr. M. Brech-biihl (ETHZ) supplied total chemical analyses of the soil sam-ples. The first author is grateful to Dr. D. Berggren, Dr. J.P.Gustafsson and M. Simonsson at SLU for many valuable dis-cussions as well as for introduction to FIA-, IR-, and TEM-analyses. Dr. V.C. Farmer provided valuable comments onan earlier manuscript. Gratitude is also expressed to threeanonymous reviewers for their constructive criticism. Thisstudy was supported by grants from the Swiss Federal Officeof Environment, Forest and Landscape, the Swiss NationalScience Foundation (grant no. 8220-046488) and the SwedishUniversity of Agricultural Sciences, Uppsala.


studies, and modelling. Ph.D. dissertation, Univ. Berne, Swit-zerland.

FAO. 1988. Soil map of the world. Revised legend. Food and Agricul-ture Organization of the United Nations, Rome.

Farmer, V.C. 1974. The Infrared Spectra of Minerals. MineralogicalSociety, London.

Farmer, V.C. 1987. The role of inorganic species in the transport ofaluminium in podzols. p. 187-194. In D. Righi and A. Chauvel(ed.) Podzols et podzolization. AFES et INRA, Paris.

Farmer, V.C, and A.R. Fraser. 1982. Chemical and colloidal stabilityof soils in the A^Oj-FeiOj-SiCVHzO system: their role in podzoliza-tion. J. Soil Sci. 33:737-742.

Farmer, V.C., A.R. Fraser, and J.M. Tait. 1979. Characterization ofthe chemical structures of natural and synthetic aluminosilicategels and sols by infrared spectroscopy. Geochim. Cosmochim.Acta 43:1417-1420.

Farmer, V.C., J.D. Russel, and M.L. Berrow. 1980. Imogolite andproto-imogolite allophane in spodic horizons: Evidence for a mo-bile aluminum silicate complex in podzol formation. J. Soil Sci.31:673-684.

Farmer, V.C., J.B. Russell, and B.F.L. Smith. 1983. Extraction ofinorganic forms of translocated Al, Fe and Si from a podzol Bshorizon. J. Soil Sci. 34:571-576.

Fordham, A.W., and K. Norrish. 1983. The nature of soil particlesparticularly those reacting with arsenate in a series of chemicallytreated samples. Aust. J. Soil Res. 21:455-477.

Foy, C.D. 1984. Physiological effects of hydrogen, aluminum andmanganese toxicities in acid soil. p. 57-97. In F. Adams (ed.) Soilacidity and liming. 2nd ed. ASA, Madison, WI.

Gee, G.W., and J.W. Bauder. 1986. Particle-size analysis, p. 383-411.In A.L. Page et al. (ed.) Methods of Soil Analysis. Part 1. Physicaland Mineralogical Methods. 2nd ed. ASA, Madison, WI.

Gustafsson, J.P., P. Bhattacharya, D.C. Bain, A.R. Fraser, and W.J.McHardy. 1995. Podzolisation mechanisms and the synthesis ofimogolite in northern Scandinavia. Geoderma 66:167-184.

Gustafsson, J.P., D.G. Lumsdon, and M. Simonsson. 1998. Aluminiumsolubility characteristics of spodic B horizons containing imogolite.Clay Minerals 33:77-86.

Kendall, M.G., and A. Stuart. 1967. The advanced theory of statistics.Volume 2. Inference and relationship. 2nd edition. Griffin, London.

Kittrick, J.A. 1969. Soil minerals in the Al2O3-SiO2-H2O system anda theory of their formation. Clays Clay Miner. 17:157-167.

Kittrick, J.A. 1970. Precipitation of kaolinite at 25°C and 1 atm. Claysand Clay Miner. 18:261-267.

Koch, O.G., and G.A. Koch 1964. Handbuch der Spurenanalyse.Springer-Verlag, Heidelberg, Germany.

LaZerte, B.D., and J. Findeis. 1995. The relative importance of oxalateand pyrophosphate extractable aluminum to the acidic leachingof aluminum in Podzol B horizons from the Precambrian Shield,Ontario, Canada. Can. J. Soil. Sci. 75:43-54.

Lumsdon, D.G., and V.C. Farmer. 1995. Solubility characteristics ofproto-imogolite sols: how silicic acid can de-toxify aluminium solu-tions. Eur. J. Soil Sci. 46:179-186.

Luster, J., A. Yang, and G. Sposito. 1993. On the interpretation oflabile aluminum as determined by reaction with 8-hydroxyquino-line. Soil Sci Soc. Am. J. 57:976-980.

Manley, E.P., W. Chesworth, and L.J. Evans. 1987. The solution chem-istry of podzolic soils from the eastern Canadian shield: a thermody-namic interpretation of the mineral phases controlling soluble Al3+and H4SiO4. J. Soil Sci. 38:39-51.

Matzner, E. 1992. Acidification of forests and forest soils: Currentstatus. Studies in environmental science 50:77-86.

Mulder, J., and A. Stein. 1994. The solubility of aluminum in acidicforest soils: Long-term changes due to acid deposition. Geochim.Cosmochim. Acta 58:85-94.

Mulder, J., N. van Breemen, and H.C. Eijck. 1989. Depletion of soilaluminium by acid deposition and implications for acid neutraliza-tion. Nature 337:247-249.

Neal, C, R.A. Skeffington, R. Williams, and D.J. Roberts. 1987. Alu-minium solubility controls in acid waters: the need for a reappraisal.Earth Planet. Sci. Lett. 86:105-112.

Nordstrom, D.K., and H.M. May. 1996. Aqueous equilibrium datafor mononuclear Al species, p. 39-80. In G. Sposito (ed.) Theenvironmental chemistry of aluminum. 2nd. ed. Lewis Publ., BocaRaton, FL.

Paces, T. 1985. Sources of acidification in Central Europe estimatedfrom elemental budgets in small basins. Nature 315:31-36.

Palmer, D., and D.J. Wesolowski. 1992. Aluminum speciation andequilibria in aqueous solution: II. The solubility of gibbsite in acidicsodium chloride solutions from 30 to 70°C. Geochim. Cosmochim.Acta 56:1093-1111.

Parfitt, R.L., and C.W. Childs. 1988. Estimation of forms of Fe andAl: A review, and analyses of contrasting soils by dissolution andMoessbauer methods. Aust. J. Soil Res. 26:121-144.

Parfitt, R.L., and J.M. Kimble. 1989. Conditions for formations ofallophane in soils. Soil Sci. Soc. Am. J. 53:971-977.

Pokrovski, G.S., J. Schott, J.-C. Harrichoury, and A.S. Sergeyev. 1996.The stability of aluminum silicate complexes in acidic solutionsfrom 25 to 150°C. Geochim. Cosmochim. Acta 60:2495-2501.

Ross, G.J., and H. Kodama. 1979. Evidence for imogolite in Canadiansoils. Clays Clay Miner. 27:297-300.

Russel, J.D., and A.R. Fraser. 1994. Infrared methods, p. 11-67. InM.J. Wilson (ed.) Clay mineralogy: spectroscopic and chemicaldeterminative methods. Chapman and Hall.

Schecher, W.D., and C.T. Driscoll. 1987. An evaluation of uncertaintyassociated with aluminum equilibrium calculations. Water Resour.Res. 23:525-534.

Schwertmann, U. 1964. Differenzierung der Eisenoxide des Bodensdurch Extraktion mil Ammoniumoxalat-Losung. Z. Pflanzenern.,Dong., Bodenk. 105:194-202.

Soil Survey Staff. 1996. Keys to soil taxonomy. 7th ed. U.S. Print.Office. Washington, DC.

Stumm, W., and J.J. Morgan. 1981. Aquatic chemistry. Wiley Intersci-ence, New York.

Su, C., and J. Harsh. 1994. Gibbs free energies of formation at 298Kfor imogolite and gibbsite from solubility measurements. Geochim.Cosmochim. Acta 58:1667-1677.

Su, C, J.B. Harsh, and J.S. Boyle. 1995. Solubility of hydroxy-alumi-num interlayers and imogolite in a spodosol. Soil Sci. Soc. Am.J. 59:373-379.

Sumner, M.E., M.V. Fey, and A.D. Noble. 1991. Nutrient status andtoxicity problems in acid soils, p. 149-182. In B. Ulrich and M.E.Sumner (ed.) Soil acidity. Springer, Berlin.

Sverdrup, H., P. Warfvinge, L. Blake, and K. Goulding. 1995. Model-ling recent and historic soil data from the Rothamstead experimen-tal station, England, using SAFE. Agriculture Ecosystems and En-vironment 53:161-177.

Tait, J.M., N. Yoshinaga, and B.D Mitchell. 1978. The occurrence ofimogolite in some Scottish soils. Soil Sci. Plant Nutr. (Tokyo)24:145-151.

Thomsen, J., K. Johnson, and R. Petty. 1983. Determination of reactivesilicate in seawater by flow injection analyses. Anal. Chem. 55:2378-2382.

Turner, R.C., and J.E. Brydon. 1965. Factors affecting the solubilityof A1(OH)3 precipitated in the presence of montmorillonite. SoilSci. 100:176-181.

Turner, R.C., and J.E. Brydon. 1967. Effect of length of time on someproperties of suspensions of Arizona bentonite, illite, and kaolinitein which aluminum hydroxide is precipitated. Soil Sci. 103:111-117.

Ulrich, B. 1989. Effects of acidic precipitation on forest ecosystemsin Europe. In D.C. Adriano and A.H. Johnson (ed.) Acidic precipi-tation. Volume 2: Biological and ecological effects. Springer,New York.

Van Breemen, N., C.T. Driscoll, and J. Mulder. 1984. Acid depositionand internal proton sources in acidification of soils and waters.Nature 307:599-604.

Walker, W.J., C.S. Cronan, and P.R. Bloom. 1990. Aluminum solubil-ity in organic soil horizons from northern and southern forestedwatersheds. Soil Sci. Soc. Am. J. 54:369-374.

Wang, C., J.A. McKeague, and H. Kodama. 1986. Pedogenic imogoliteand soil environments: case study of spodosols in Quebec, Canada.Soil Sci. Soc. Am. J. 50:711-718.

Webster, R. 1997. Regression and functional relations. Europ. J. SoilSci. 48:557-566.

Xu, S., and J.B. Harsh. 1995. Influence of autocorrelation and mea-surement errors on interpretation of solubility diagrams. Soil Sci.Soc. Am. J. 59:1549-1557.

Yuan, T.L. 1959. Determination of exchangeable hydrogen in soilsby titration method. Soil Sci. 88:164-167.

ERRATUMAluminum Solubility Control in Different Horizons of a Podzol

M. ZYSSET, P. BLASER, J. LUSTER, AND A.U. GEHRINGSoil Sci. Soc. Am. J. 63:1106-1115 (September-October 1999).

Errors were found in a chemical formula in the paragraph immediately above the Conclusions section. Thestatements in question on page 1114 should readKittrick (1970) pointed out that stability of kaolinite in soils might vary between the crystalline form and that ofhalloysite. Thus, log *KS for 0.5 Al2Si2O5 (OH)4 + 3H+ = A13+ + H4SiO4 + 0.5 H2O might vary between 3.1 and5.4 at 20°C (thermodynamic data reported by Manley et al, 1987; Stumm and Morgan, 1981; and Kittrick, 1969).

aluminum solubility control in different horizons of a podzol

Documents