The role of organic matter in controlling aluminum solubility in acidic mineral soil horizons

Geochimica et Cosmochimica Acta, Vol. 59, No. 20, pp. 4167-4180, 1995 Copyright 0 1995 Elsevier Science Ltd Printed in the USA. All rights reserved 0016-7037/95 $9.50 + .OO

Pergamon

0016-7037(95)00279-O

The role of organic matter in controlling aluminum soluhility in acidic mineral soil horizons Department

DAN BERGGREN*.’ and JAN MULDER’ of Soil Science and Geology, Wageningen Agricultural University, PO Box 37, 6700 AA Wageningen, The Netherlands (Received

August 5, 1994; accepted in revised form June 15, 1995)

Abstract-Despite the ecological importance of potentially phytotoxic Al, its solubility control in acidic mineral soils remains unresolved. We examined the solpbility of Al in mineral horizons of two acidic forest soils (Inceptisol and Spodosol) in southern Sweden using a series of batch experiments. Dissolution of Al was found to consist of a rapid solubilization of reactive solid phase Al, which quickly reached an equilibrium state, superimposed on a slow dissolution of less reactive Al-containing phases (e.g., primary Alsilicates). Titration experiments in the pH range 3.2-4.7 using an equilibration time of 5 days showed that at pH < 4.1, all suspensions were undersaturated with respect to gibbsite (Al (OH), ; log * KS0 = 8.85 at 8°C). Under such conditions, the Al solubility could be explained qualitatively by equilibrium complexation reactions with soil organic matter. Quantitatively, our results could be reproduced reasonably well using the mechanistic model WHAM, which describes the binding of Al by humic substances in organic soils. This suggests that the pool of organically bound soil Al controls the Al solubility in suspensions of strongly acidic soils. Due to the kinetically constrained release of Al from primary and secondary minerals, the amount of organically bound Al, and therefore the Al solubility in the suspensions, gradually increases with time. Consequently, a quantitative evaluation of Al solubility data from long-term batch experiments should consider both equilibrium and kinetic processes. 1. INTRODUCTION

summer, when the residence time of soil water is longest and temperatures are highest (Mulder and Stein, 1994). Recently, Mulder et al. (1989b) proposed that the Al”+ activity in strongly acidified mineral soil horizons is regulated by complexation reactions with soil organic matter. These investigators also found that the pyrophosphate-extractable pool of Al (a measure of the organically bound Al pool) is the most reactive Al fraction during leaching with a dilute strong acid. Upon depletion of this pool, leaching solutions became increasingly undersaturated with respect to Al(OH)X. Dahlgren and Walker ( 1993) confirmed the fast dissolution kinetics of organically complexed solid-phase Al in a study of three Spodosol Bs horizons. Nevertheless, Dahlgren and Walker ( 1993 ) suggested that at equilibrium the activity of Al ‘+ was controlled by interlayer hydroxy-Al, without attempting to interpret this apparent contradiction. Previous studies performed with organic soil layers (Cronan et al., 1986; Walker et al., 1990) have shown that dissolved Al rapidly attains equilibrium with the solid phase. However, for mineral soil layers few tests exist to determine whether dissolved Al attains a rapid equilibrium with resident soil organic matter. In addition, the conditions under which the activity of Al’+ in solution is controlled by organic complexation reactions, as opposed to being controlled by the dissolution of Al ( OH)S, remain unclear. The objectives of the present study were to determine, both qualitatively and quantitatively, if the solubility of Al at low pH in some selected mineral horizons can be explained by equilibrium binding to soil organics, and to elucidate the conditions under which this equilibrium mechanism is in operation.

Because dissolution of aluminum from the soil solid-phase is the dominant mechanism of proton buffering in many forest soils of Northern Europe (Matzner, 1989; Mulder et al., 1989a; Berggren, 1992), it is a key process in soil acidification models. Despite the importance of Al in acid neutralization in acidic mineral soils, the processes controlling the solution activity of Al’+ are not fully understood. Published acidification models (e.g., Birkenes, MAGIC, SAFE, SMART) have generally assumed that Al’+ activity is regulated by equilibrium with an Al ( OH)3 phase. Some studies of the solution chemistry in Spodosol B horizons support this concept (David and Driscoll, 1984; Dahlgren et al., 1989), while others do not. In particular, soil solutions with pH < 4.1 are generally undersaturated with respect to Al ( OH), (Matzner, 1989; Mulder et al., 1989a; Berggren, 1992; Mulder and Stein, 1994). In general, the degree of undersaturation seems to increase with decreasing pH (Seip et al., 1989; Berggren, 1992; Matzner and Prenzel, 1992; Mulder and Stein, 1994). Kinetically constrained mineral dissolution has been proposed as a possible explanation for the apparent undersaturation at low pH values (e.g., Matzner and Prenzel, 1992). However, in a controlled laboratory experiment, Spodosol Bs horizons readily attained equilibrium with gibbsite (May et al., 1979), starting from conditions of both oversaturation and undersaturation (Dahlgren et al., 1989). Furthermore, field observations indicate that undersaturation reaches its maximum in

* Present address: Department of Soil Sciences, Swedish University of Agricultural Sciences, P.O. Box 7014, 750 07 Uppsala, Sweden. + Author to whom correspondence should be addressed. * Presenr address: Norweigian Institute of Forest Research, skoleveien 12, 1432 As, Norway.

2. THEORY Equilibrium of a soil solution with an Al( OHX phase can be expressed as

Heg4167

4168

D. Berggren and J. Mulder

Table 1. Thermodynamic data applied in chemical equilibrium calculations using ALCHEMI. Reaction

log

K25 oc

where x is the stoichiometry of the reaction. The equilibrium equation for this reaction may be formulated as (Al”+} [H,R]/( [AIR”-“‘] (H+)“)

AH,(kJ/mol)

Al3+ + H20 = AI(OH)z+ + H+

-4.958

54.8”

Al3+ + 2H20 = AI(OH)2+ + 2H+

-10.6a

1oF

Al3+ + H&i04

-2.5OC

0”

= K,

(5)

where squared brackets represent the concentration (in, e.g., mol/kg SOM) of the complexation sites and where K is an apparent equilibrium constant. In a logarithmic form Eqn. 5 may be written as pA1 = x pH - log (K[AlR’“-“‘]I[HJ?]).

= AIH$iO$+

+ H+

AI + 3H+ = Al3+ + 3Hzo gibbsite

7.74e

In contrast to the equilibrium with an Al(OH)3 phase, the equilibrium of Al with SOM results in a pH-pAl plot with a variable slope. The slope varies depending on the type of sites involved in the binding of Al by SOM. Also, the value of the intercept, which is a function of K and of the saturation of the organic complexation sites with Al, is variable.

-105e

aWESOLOWSKl and PALMER (1994). bcalculated from data given in WESOLOWSKI and PALMER (1994). ~FARMER and LUMSDON (1994). dAH,was set to zero because no value was found in the literature. ~PALMER and WESOLOWSKI 1992.

Al(OHh(s)

+ 3H’ @ Al”

Table 2. Selected characteristics of the c2-mm size fraction of the Skanes V&sjo (SkV) and Rosenlund (Ros) soils. Also given are optimized values of humic acids (CHA), fulvic acids (CFA), and reactive Al (CAI) for the SkV Bh sample, as obtained using the mechanistic model WHAM (TIPPING, 1994).

+ 3H20.

This equilibrium results in a cubic relationship between Al ‘+ and H + according to (Al’+}/( H+}3 = *KS,,,

(1)

where curled brackets indicate activities and where *KS0 is the solubility product constant of Al (OH), . In a logarithmic form, Eqn. 1 becomes pA1 = 3 pH - log * Kso.

(2)

Equation 2 indicates that equilibrium with Al( OH)3 results in a pH-pA1 plot with a slope of 3 and an intercept log * Kso. *KS, depends on the crystallinity and purity of the Al( OH), phase. For example, the most frequently used log *KS, value for pure macrocrystalline gibbsite is 8.11 at 25°C (May et al., 1979). Recently, Palmer and Wesolowski (1992) revised this log *KS, value to 7.74, amongst others based on the May et al. ( 1979) experimental data, but accounting for Al complexation by organic ligands in pH buffers. For a given solution, the ion activity product (Q) for AI may be calculated according to Q = (Al’+}/(H+}‘.

(6)

(3)

SkV Bh

SkV Bs

ROSA

Ros B

0.10 0.05 0.04 0.04 3.2 0.53

0.06 0.02 0.02 0.02 1.8 0.04

0.16 0.04 0.03 0.02 2.0 0.10

0.21 0.02 0.02 0.01 0.80 0.01

5.3 8.7 15 350

25 10 20 1300

8.1 6.5 8.3 970

12 4.0 4.2 2800

Fe

5.3 10

31 14

8.3 8.5

13 4.7

a ate-dithionite Extrm Al ;cmol/kg) Fe

5.7 16

29 19

8.2 16

11 11

moer chloride Extra Al (cmol/kg)

2.1

5.6

2.4

1.2

Carbon (g/kg)

24

28

19

6.8

Ootimized Dara eterS CHAp;9/kg)b CFA (gfig)b CAI (cmol/kg)

8.1 6.6 1.5

Particle Size % Sand % Silt % Clay

71.5 25 3.5

66.5 31 2.5

65 28 7

64.5 31 4.5

Property Exchanaeable Ca (cmol&g)a Mg Na K Al H Pvroohosohate Al (cmol/kg) Fe (cmol/kg) C (g/kg) cmol Al/kg C

cations

Extract

Furthermore, we define SI = log (Q/*Kso),

(4)

where SI is the saturation index, which is used to express the degree of saturation of the solution with respect to the Al ( OH)s phase under consideration; SI > 0 indicates oversaturation, SI < 0 undersaturation, and SI = 0 saturation with respect to the mineral phase considered. Complexation of Al by soil organic matter (SOM) may schematically be expressed as (cf. Tipping and Hurley, 1992) AU?‘“-“‘+ + xH+ G+Al’+ + HJ?, where R indicates a complexation site located on SOM and

acmol&g = cmol charge equivalent per kg dry soil. bOne g FA or HA is approximately equal to 0.5 g organic C.

Solubility of Al in soils containing organic matter Table 3. Clay mineralogy of the c2-urn size fraction of the Skanes Varsjd (SkV) and Rosenlund (Ros) soils. +++ = major component, ++ = minor, + = trace, - = not detected

Hydroxy-Al interlayered vermiculite Vermiculite Kaolinite Chlorite Mica aData are from

OLSSON

3.

SkVBh

SW Bsa

ROSA

Ros B

+++ ++ ++

+++ ++ ++

+++ + ++ + +

+++ + ++ + +

and

( 1989). Samples used in the present study were taken from the Bh (denoted B/A by Olsson and Melkerud, 1989) and Bsl horizon at Sk&es Vi+rsjSand from the A and B horizon at Rosenlund. At both sites Norway spruce (P&a dies (L.) Karst.) is the predominant tree species. Ground vegetation at Sties Varsjo is dominated by the grass Deschumpsiu jlexuosa (L.) with patches of the dwarf shrub Vaccinium myrfilh (L.) and mosses (Bergkvist, 1987). At Rosenlund, mosses occur, but no undergrowth of higher plants is present. At both sites the climate is maritime, with precipitation relatively evenly distributed over the year. The mean annual precipitation is 795 mm, and the monthly mean temperature ranges from 16.6”C in July to - 1.5”C in January and February. 3.2. Soil Characterization

MELKERUD (1989).

MATERIAL AND METHODS

3.1. Study Sites The two study sites, Skanes Wirsjii (56”40’N, 13”30’E) and Rosenlund (56”05’N, 13”55’E), are located in southernmost Sweden. The soil is a Typic Haplorthod at Sk%nes Varsjb (SkV) and an Eutric Haplumbrept at Rosenlund (Ros) (SOIL SURVEY STAFF, 1992). Both soils developed in sandy loamy glacial till, mainly originating from siliceous rock (gneiss and granite). At Rosenlund, the parent material also contains fragments of limestone. A detailed description of the soil profile at Skt!nes V%rsjGis given by Olsson and Melkerud

SkV Bh

4169

Field-moist soil samples were passed through a 2 mm plastic sieve to remove coarse fragments and to homogenize the samples. Dry weight was determined at 105°C and organic carbon was determined using a LECO CR 12 carbon analyzer. All extractions were performed on field moist samples using an end-over-end shaker. Solidphase Al was determined by separate extractions with ( 1) 0.1 M Ba(N02)* for 2 h, (2) 0.5 M CuCl, for 2 h (Jou and Kamprath, 1979), (3) Na,P,O, for 16 h at pH 10 (McKeague, 1967), (4) acid NH, oxalate for 4 h at pH 3 in the dark ( McKeague and Day, 1966), and ( 5) citrate-dithionite for 16 h (Holmgren, 1967). Clear solutions were obtained by centrifugation, followed by filtration through 0.2 pm membrane filters. Exchangeable amounts of Ca’ +, Mg *+, Na +, and K+ were estimated by measuring their concentration in the Ba(NO?)* extracts. A measure of exchangeable H’ was calculated based on the pH and the total Al concentration of the Ba(N03b

0 Sum 8C-Na ? ?H,SiO,

0 •I A o

0.0

10.0

20.0

30.0

Time (h’“)

Time (h”*) (c>

10.0

Time (h”‘)

1

-2.06

8/ 0.0 -

Sum SC H,SIO, H+ Al” _

20.0

2.0

I

.

I

0 Sum 8C-Ca ? ?H.SiO,

Ros B

10.0

20.0

Time (h’2>

FIG. 1. Concentrations of major solutes as a function of the square root of the reaction time (Experiment 1). Two g of field moist soil samples were. equilibrated in 10 mL of a solution having a concentration of 1.5 mM HNO-, and 3.5 mM NaNOS (SW Bh and Ros A) or 5 n&l HN07 (SkV Bs and Ros B). BC denotes the sum of Na+, K’, Ca”, and Mg2’. The values given at t = 120 h and t = 719 h are saturation indices for gibbsite (log *KS0 = 8.85 at 8°C).

- 64 .

D. Berggren and J. Mulder

4170

Table 4. Concentrations of H+, OH-, Na+, and NOs- (corrected for dilution with soil water resident in the field moist samples) in solutions added to soil samples in Experiment 2. Also give are the concentrations of total Al (Al!,t) and quickly reacting Al (Alqr) and pH in the soil suspensions after 5 d of equilibration.

Added

[H+l

W-l

W+l

___

Final ___

PJOs-1

b%tla

[Alqrlb

PH

mM

0 0 0 0.194 0.388 0.582 0.775 1.07 1.45 1.94

0.388 0.194 0 0 0 0 0 0 0 0

5.74 5.78 5.82 5.58 5.35 5.12 4.88 4.54 3.39 2.91

SkV Bh 5.35 5.58 5.82 5.78 5.74 5.7 5.66 5.6 4.85 4.85

NDC ND ND ND ND 0.193 0.196 0.248 0.330 0.434

0.026 0.029 0.039 0.048 0.068 0.089 0.109 0.155 ND ND

4.04 3.94 3.85 3.76 3.68 3.59 3.53 3.46 3.35 3.27

0 0 0.482 0.964

0.193 0 0 0 0 0 0 0 0 0 0

9.44 9.64 8.87 7.71 6.75 5.78 3.86 1.93 0 0 0

P 9.64 9.16 8.67 8.19 7.71 6.75 5.78 4.82 5.78 9.64

ND ND 0.269 0.421 0.584 0.747 1.08 1.42 1.65 1.98 3.26

0.082 0.091 0.195 0.33 0.481 0.606 0.96 1.32 1.83 ND ND

4.61 4.58 4.44 4.34 4.24 4.18 4.08 3.98 3.91 3.85 3.68

0.388 0.194 0 0 0

5.74 5.78 5.82 5.59 5.36 5.12 4.89 3.78 3.4 2.62

&XU 5.36 5.59 5.82 5.78 5.74 5.7 5.67 4.85 4.85 4.56

ND ND ND ND ND 0.204 0.245 0.328 0.441 0.562

0.012 0.019 0.037 0.061 0.081 0.13 0.171 0.235 0.344 0.453

4.69 4.54 4.41 4.28 4.19 4.09 4.03 4.01 3.9 3.82

9.55 9.74 8.77 7.79 7.5 6.43 4.29 2.14 1.46 0

&XI 9.35 9.74 9.26 8.77 8.96 8.38 7.21 6.04 6.33 5.85

ND ND 0.169 0.315 0.489 0.639 0.912 1.092 1.275 1.368

0.015 0.03 0.107 0.233 0.359 0.502 0.743 0.943 1.143 1.286

4.89 4.79 4.58 4.45 4.35 4.26 4.13 4.02 3.91 3.84

1.45 1.93 2.89 3.86 4.82 5.78 9.64

0 0 0 0.194 0.388 0.582 0.776 1.07 1.46 1.94

: 0.487 0.974 1.46 1.95 2.92 3.9 4.87 5.85

: 0 0 0

0.195 0 0 0 : 0 0 0 0

aConcentration determined bconcentration determined CND = not determined.

by AAS. according to CLARKE et al. (1992).

extracts according to Meiwes et al. (1984). Extracted Al, Fe, Ca, Mg, Na, and K were determined by atomic absorption spectroscopy (AAS). Dissolved organic C was determined in the Na.+P20, extracts using a Shimadzu Total Organic Carbon Analyzer TOC-500. Clay mineralogy was determined using X-ray diffraction analysis.

3.3. Analytical

Procedures

Speciation of Al in solutions was performed using a recently developed FL4 (Flow Injection Analysis) method (Clarke et al., 1992) based on the principle that Al species differ in the rate at which they

Solubility of Al in soils containing organic matter

4171

5.0 4.8

4.5

s

4.0

3.5 3.2 I 3.0

I 3.5

1 4.0

I 4.5

3.0 3.6

3.8

4.0

PH 6.01

’

’

I

1

’

I

4.4

4.6

4.8

PH ’

I

’

I

Ros A

1

4.2

1 (c) a

5.0

lh ? ? 5h 0 Id oo 5d v l5d A 30d

Ros B

04

% 4.0

3.0

4.0

-

3.6

3.8

4.0

4.2

4.4

4.6

4.8

3.0 I3.0

4.0

PH

PH

FIG. 2. pA1 vs. pH for suspensions of (a) SkV Bh, (b) SkV Bs, (c) Ros A, and (d) Ros B samples. Data from Experiment 1 are represented by open symbols; data from Experiment 2 by filled circles. In both experiments two g of field moist soil samples were equilibrated in 10 mL of solution. For Experiment 2 the chemical composition of the added solutions as well as pH and concentrations of Al,, and AI,, in the equilibrated suspensions are given in Table 4. For the two highest acid additions to SkV Bh and SW Bs we assumed Ah, to be 80% (SkV Bh) and 100% (SW Bs) of Al,,,. The solid line indicates the solubility of gibbsite (log *KS0 = 8 85 at 8’C), and the arrow indicates the suspension with no acid or base added. react with I-hydroxoquinoline. The method has been thoroughly tested on model systems (Clarke et al., 1991, 1992) as well as on natural waters (Berd6n et al., 1994). The measured Al fraction, denoted as quickly reacting Al (Als,), includes Al”, AlOH*‘, AI(O AlSOf , and AlH$iiO, ‘+ . The analytical conditions given by Clarke et al. (1992) could only be used for Al,, concentrations less than circa 0.2 mM. However, the analytical range could be increased substantially by decreasing the injection volume and carrier stream flow rate. For example, using an injection volume of 10 /JL and a carrier stream flow rate of 0.16 mL/min, Alsr concentrations of 2 mM could be measured. Concentrations of All+ were calculated using the chemical equilibrium program ALCHEMI (Schecher and Driscoll, 1987) version 4.0, considering the species Al”, A10H2+, Al( OH):, and AlH,SiO:’ The formation constants adopted in these calculations are given in Table 1. The total concentration of Al was determined by AAS. The pH was measured potentiometrically with a combined pH electrode (GK2421C, Radiometer, Denmark) calibrated at pH 4 and 7 using Orion low ionic strength buffers. Dissolved Si was measured using the blue silicomolybdic acid procedure (Weaver et al., 1968) (Experiments 1, 2) or according to Velthorst ( 1993) (Experiment 4). Total concentrations of dissolved Al, Ca, Mg, Mn, Na, and K were determined by AAS. For the determination of Ca and Mg we ad&d 72 mM LaCl? and for Na and K 7.5 mM CsCl to standard and sample solutions.

starting the experiments. However, in Experiment 4 the samples were thawed and then stored in a refrigerator at circa 6°C 11 or 14 weeks prior to the experiment. In all experiments 10 mL of solution was ad&d to 2.0 g field-moist soil, equivalent to 1.6 (SkV Bs) and 1.7 g dry soil (SkV Bh, Ros A and Ros B), in 15 mL polypropylene centrifuge tubes with rubber-lined Teflon screw caps. Centrifuge tubes were placed horizontally in a thermostated shaking water bath (Hetomix TBVS 01/02 equipped with cooling bath Hetofrig CB60VS; Heto Lab Equipment A/S, Denmark) and shaken at circa 165 rpm to avoid sedimentation of soil particles. A temperature of 8.0 ? O.l”C was used in all experiments. After equilibration, the suspensions were cenbifuged at 8°C. The pH of supematants was measured immediately after centrifugation in a cold room, also at about 8°C. Another portion of the supematants was filtered through a 0.2 pm single-use filter (Acrodisc PF, Gelman Sciences, MI) prior to analysis of Si, AG,, and base cations. In all experiments, the concentration of NO;, which was the only inorganic anion present in significant concentrations, was calculated from the added amount of NO:, corrected for dilution with soil water resident in the field-moist samples.

3.4. Bat& Experiments-General The field-moist, sieved samples were kept frozen until further use. In Experiments 1, 2, and 3, samples were thawed immediately before

Aluminum was the dominant exchangeable cation in all four soil horizons (Table 2). In general, pyrophosphate (Na&O,), oxalate, and citrate-dithionite extracted about

4. RESULTS 4.1. Soil Characteristics

D. Berggren and J. Mulder

4172

Table 5. Estimated increase in the amounts of reactive Al between days 5 and 31 for 2.0 g of field-moist soil suspended in 10 ml solution of 1.5 mM HNOo and 3.5 mM NaN03 (SkV Bh and Ros A) or 5 mM HNOs (SkV Bs and Ros B). The estimates were obtained through a CuClp extraction procedure (JOU and KAMPRATH,1979). Means of duplicate analyses are given, and values in parentheses are the corresponding ranges. Also shown are calculated amounts of Al released to the suspensions through silicate weathering during the same period (see text for assumptions and calculations).

ogical composition between the soils. All soils mainly consisted of quartz (more than 70%) and feldspar-s with minor amounts of amphibole. The content of plagioclase exceeded that of K-feldspar. The clay fraction of the four soils had a similar composition, with hydroxy-Al interlayered vermiculite as the major component and with minor amounts of vermiculite and kaolinite (Table 3). 4.2. Batch Experiments

Sample

CuCls extractable Al (umof)

SkV Bh SkV Bs Ros A Ros B

4.1 (1.2) 13 (10) 6.1 (0.7) 11 (3.5)

Al release from silicate weathering (umol) 2.2 1.4 1.9 3.2

equal amounts of Al. Only for the SkV Bs sample N%P207 was somewhat less effective than the other extractants, removing about 80% of the Al extracted by oxalate and citratedithionite. The efficiency of N&P20, in extracting Al relative to that of oxalate and citrate-dithionite suggests that the major form of noncrystalline Al in all horizons was bound to SOM, and that short-range-ordered Al silicates like allophane and imogolite were not present in any significant amounts (Par&t and Childs, 1988; Dahlgren and Walker, 1993). The C content was low in all horizons, ranging from 6.8 g/kg in Ros B to 28 g/kg in SkV Bs. The ratio of Al to C in the N&P207 extracts ranged from 350-2800 cmol/kg organic C. If NhP207 extracts only organically bound Al, then the Al saturation of the organic matter in the four soil horizons would increase in the order: SkV Bh < Ros A < SkV Bs < Ros B. However, after comparing Al + Fe to C molar ratios in N%PZ07 extracts of Spodosol B horizons with the number of potential binding sites per mol C, Higashi et al. ( 1981) concluded that NQ207 may also extract some Fe and Al not bound to organic matter. On the basis of published values of the COOH content of SOM, these authors proposed 1000 cmol/kg C as a maximum value for metal binding by SOM. A recent investigation on metal binding by FA and HA (Tipping and Hurley, 1992) suggests that this maximum value is reasonable. Thus, the Al to C ratios given in Table 2 suggest that for the SkV Bs and Ros B samples NhP207 also extracted some inorganic forms of Al. This N&P207-extractable inorganic Al fraction may include amorphous surface precipitates of Al ( OH)3 ( McKeague et al., 197 1) , coprecipitates of amorphous Al ( OHb and humic substances, sorbed polynuclear Al species, and interlayer hydroxy-Al. Copper chloride has been proposed as an alternative extractant of organically bound Al (Hargrove and Thomas 1981; Walker et al., 1990). The amount of Al extracted by CuCl, was significantly less than the amount extracted by N&P207 (Table 2). The ratio of CuCl*-extractable Al to N&P,O,-extractable C ranges from 140-290 cmol/kg. This stays well below the maximum value as indicated above suggesting a higher degree of specificity of CuC& to extract mononuclear Al bound to SOM. X-ray diffraction analysis of the lo-50 pm fraction showed that there were no significant differences in mineral-

4.2.1. Experiment

1: Kinetics

Kinetics of the mobilization of major compounds from the solid phase were studied in duplicate. Samples were equilibrated for 1 h, 5 h, 24 h, 5 days, 15 days, and 30 days in a solution of 1.5 mM HN03 and 3.5 mM NaNO? (SkV Bh and Ros A) or 5 mM HNO, ( SkV Bs and Ros B ) . Solutions were analyzed for pH, Si, Al,,, Ca, Mg, Mn, Na, and K. The experimental results, which were highly reproducible, indicate that a major part of the added H+ was neutralized by a rapid release of Al and base cations from the solid phase in all samples (Fig. 1) . The maximum concentration of Al was reached within one hour. The concentration of Si was initially low, suggesting a lack of quickly reacting Al silicates in these samples, which is in accordance with the observed distribution of extractable Al discussed above. Subsequently, the Al”+ concentration decreased slowly. This decrease was accompanied by a decline in the concentration of H+ due to the slow weathering of primary minerals, as indicated by the simultaneous increase of base cations and Si. In the Ros B sample a substantial release of Ca*’ occurred during the first 24 h (Fig. Id). This Ca*+ may originate from traces of CaCO? preserved in the B horizon. Van Grinsven (1988) also reported the release of Ca*+ from an acidified, previously calcareous soil upon exposure to a strong acid. This author proposed that Ca-containing minerals, previously stabilized in the acid environment by coatings of amorphous Al (OH), , may solubilize after dissolution of the coatings by strong mineral acids. As indicated in Fig. 1, the solutions in our experiment were all undersaturated with respect to macrocrystalline gibbsite after 5 days of equilibration, the SkV Bh and Ros A samples being more undersaturated than the SkV Bs and Ros B samples. Notwithstanding the obvious undersaturation with respect to gibbsite, the A13+ activity decreased between days 5 and 30, suggesting that the activity of A17’ is not controlled by a kinetically constrained dissolution of this mineral. Apparently, other mechanisms result in an activity of Al’+ which is lower than that obtained in case of equilibrium with respect to gibbsite. Also, the observation that Al’+ and H + decreased with a stoichiometry ranging from 0. 1- 1.1 ( ApAl/ApH calculated for days 5 - 15 ) , which is significantly below the value of 3 expected for precipitation of an Al( OH)J phase, suggests that some phase other than Al (OH), regulates the activity of Al ‘+ . Based on previously published work, we propose that a major fraction of Al released from the solid phase during the first hours of the experiment originated in the organically bound Al pool (Jersak and McCall, 1989; Mulder et al., 1989b; Dahlgren and Walker, 1993). The subsequent gradual decline in H’ activity was probably the result of the weath-

Solubility of Al in soils containing organic matter ering of primary and secondary minerals releasing base cations and Al (Van Grinsven et al., 1986). The gradual decrease in the concentration of Al ‘+ can be explained by an increased (pH dependent) complexation of Al with SOM (see discussion below). The SkV Bh and Ros A suspensions remained substantially undersaturated with respect to gibbsite throughout the experiment. By contrast, the SkV Bs and Ros B suspensions became slightly oversaturated with respect to gibbsite at the end of the experiment reaching log Q (=3 pH-pA1) values of 9.08 ( SkV Bs) and 9.33 (Ros B). Also for Spodosol Bs horizons from Hubbard Brook and Bear Brook (northeastern USA), log Q values were reported to be close to 9.0 (Dahlgren et al., 1989; Dahlgren and Walker, 1993). 4.2.2. Experiment 2: Acid-base batch titrations A batch titration experiment was performed to determine the solubility of Al in the four soil samples as a function of pH. Based on results from Experiment I, we selected an equilibration time of 5 days in this experiment. Mixtures of NaOH and NaN07 or HN03 and NaNOX were added in amounts necessary to maintain a reasonably constant ionic strength (SkV Bh and Ros A: 0.006 2 0.001 M; SkV Bs: 0.009 to 0.019 M; Ros B: 0.010 % 0.001 M) for each titration curve. Because the reproducibility of the experiments was good (Fig. 1 ), batch titrations in Experiments 2 and 4 were performed on unreplicated soil samples. All solutions were analyzed for pH, Si, AJ,, and Na. In addition, Ca was determined in solutions from Ros B, where this solute dominates. Other cations did not contribute significantly to the ionic strength, as revealed by Experiment 1. Experimental conditions are given in detail in Table 4. Results from the batch titration experiment are presented as pA1 vs. pH plots in Fig. 2 (filled circles). Suspensions of the SkV Bh and Ros A samples were undersaturated with respect to gibbsite in the investigated pH range. Fitting the data with a linear regression model gave slopes of 1.5 ( SkV Bh) and 1.9 (Ros A), which are considerably less than expected from an AI( solubility control. For the SkV Bs and Ros B samples, the suspensions at pH > 4.1 were slightly oversaturated with respect to gibbsite, whereas at pH < 4.1 solutions were undersaturated. At pH > 4.4, the solubility of Al was almost identical for the SkV Bs and Ros B samples with log Q being 9.3 and the slope in the pH-pA1 plot being close to 3. This indicates the presence of a reactive Al(OHh phase with a slightly higher solubility than that of gibbsite in the SkV Bs and Ros B samples. According to Experiment I, the undersaturation with respect to this Al( OH), phase at low pH was not due to slow dissolution kinetics, but rather to depletion upon elevated additions of strong mineral acids. Note that the high Al to C ratios in Na,,P20, extracts of these samples also suggest that some inorganic secondary Al component was present in these horizons. In Fig. 2, pH-pA1 data from Experiment 1 (open symbols) are superimposed on the data from Experiment 2. A shift in both pH and pA1 was revealed by comparing 5 and 30 days of equilibration. However, the shift did not follow the stoichiometty of Al(OHh (as discussed above), nor was it in accordance with the observed Al solubility curves obtained after 5 days of equilibration. To explain the latter observation

4173

we hypothesized that there was a continuous, but kinetically constrained, release of Al from primary and secondary minerals, which gradually increased the amount of reactive Al in the suspensions. If the Al’+ activity in solutions was controlled by complexation reactions with SOM, we would expect the pH-pA1 relationship to be a function of the fraction of sites occupied by Al according to Eqn. 6; i.e., the pH-pAl relationship would be a function of the total amount of reactive Al in the suspension. If this hypothesis holds true, the observed shift in the pH-pA1 relationship resulting from the weathering-induced increase in reactive Al when going from 5-30 days of equilibration would also be expected to occur following the addition of an amount of Al salt equivalent to the amount of Al released during the 25day reaction period. The weathering-induced increase in reactive Al in the equilibrium studies was investigated in Experiment 3. In the subsequent Experiment 4, the effect of this increase on the solubility of Al was studied. 4.2.3. Experiment 3: Changes in reactive aluminum during equilibration The amount of reactive Al in a soil-solution system can be defined as dissolved Al plus the amount of solid-phase Al involved in rapid equilibrium reactions with the solution. The increase in the amount of reactive Al in equilibrated suspensions was estimated through a CuCl, extraction procedure (Jou and Kamprath, 1979). The total amount of reactive Al in the suspensions was taken to be the sum of Al in the equilibrium solution and Al in the residual solid phase extracted by 0.5 M CuCl,. Copper chloride was selected as an extractant because it is efficient in extracting organically bound Al (Hargrove and Thomas, 198 1) , which we expected to be the dominant form of reactive Al in the soil solid phase. Although inorganic polymerized hydroxy-Al may also be extracted with 0.5 M CuCl, (Jou and Kamprath, 1979), its contribution to the reactive Al in the equilibrated acid suspensions was probably small (see discussion above). Soil samples were equilibrated for 5 or 3 1 days using solutions that were the same as those used in Experiment 1. A reaction time of 31 instead of 30 days was assumed not to have a quantitatively important effect on weathering. Following equilibration, the suspensions were centrifuged and the supematants withdrawn using a Pasteur pipette. The total concentration of dissolved Al in the supematants was determined by AAS. The residual solid materials were extracted with 40 ml 0.5 M CuCl* for 2 h. After centrifugation of suspensions, supematants were passed through 0.2 pm filters and analyzed for Al by AAS. The increase in reactive Al between days 5 and 3 I, as estimated by CuC& extraction, was 4 and 6 pmol per 2.0 g fieldmoist soil for the SkV Bh and Ros A samples, whereas this was 11 and 13 pmol per 2.0 g field-moist soil for the SkV Bs and Ros B samples (Table 5). The contribution of Al from silicate weathering was calculated based on the increase in the Na, Ca, and K concentrations, assuming congruent dissolution of albite, anorthite, and K-feldspar (the predominant Al-bearing primary minerals in these soils) and assuming that all solubilized Na, Ca, and K remained in solution. These assumptions were confirmed by our Si data; i.e., the calculated increases in dissolved Si from days 5- 15 and days 15-30

4174

D. Berggren and J. Mulder Table 6. Concentrations of Al3+, H+, OH-, Na+, and NOs- (correctedfor dilution with soil water resident in the fieldmoist samples) in solutionsadded to soilsamples in Experiment 4. Also give are the concentrationsof totalAl (AIt,,,) and quicklyreactingAl (Al$ and pH in the soilsuspensions after5 d of equilibration. Run #1 and Run #2 indicateexperiments conducted at different occasions.

-Final-

]H+l

Added [OH-]

0 0 0 0 0 0.475 0.475 0.475 0.475 0.475 0.951 0.951 0.951 0.951 0.951

0 0 0.388 1.07 1.94 0 0 0 0.475 1.43 0 0 0 0 0.475

0.388 0 0 0 0 0.951 0.475 0 0 0 1.90 0.951 0.475 0 0

0 0 0 0 0 0.946 0.946 0.946 0.946 0.946 1.86 1.86 1.86 1.86 1.86

0 0.473 1.42 2.84 4.73 0 0 0.946 1.89 2.84 0 0 0 0.928 1.86

0 0 0 0

0 0 0.388 1.07 1.94 0 0 0 0.478 1.43 0 0 0 0 0.476

0.388 0 0 0 0 0.952 0.476 0 0 0 1.90 0.952 0.476 0 0

5.74 5.82 5.36 4.54 2.91 4.76 3.90 3.81 2.86 1.33 4.76 2.86 2.19 1.90 0

0 0.478 1.43 2.87 4.78

0

14.3 13.4 11.5 8.60 4.78 12.4 10.5 8.60

[Al3+]

[Na+] mM

[NOa-]

[AMa

Wqrlb

PH

Run#l

0.47: 0.476 0.476 0.476 0.476 0.952 0.952 0.952 0.952 0.952

0 0 0 0 0 0.956 0.956 0.956

0.956

5.74 5.82 5.35 4.54 2.91 4.75 3.90 3.80 2.85 1.33 4.75 2.85 2.19 1.90 0

SkVBh 5.35 5.82 5.74 5.60 4.85 5.23 4.85 5.23 4.75 4.18 5.70 4.75 4.56 4.75 3.33

0.284 0.137 0.093 0.189 0.376 0.135 0.219 0.332 0.441 0.675 0.211 0.454 0.597 0.741 0.780

0.032 0.034 0.049 0.139 0.304 0.060 0.132 0.252 0.358 0.589 0.140 0.378 0.498 0.635 0.727

4.15 3.92 3.69 3.45 3.25 3.80 3.62 3.52 3.40 3.25 3.79 3.54 3.45 3.38 3.30

14.2 0 0 13.2 0 11.3 0 8.51 0 4.73 1.89 12.3 0 10.4 0 8.51 0 6.62 0 4.73 3.71 10.2 1.86 6.50 0 4.64 0 2.78 0 0.928

SkVBs 14.2 13.7 12.8 11.3 9.46 13.2 13.2 12.3 11.3 10.4 12.1 10.2 10.2 9.28 8.35

0.111 0.228 0.530 1.02 1.67 0.315 0.945 1.28 1.63 1.90 0.582 1.27 1.87 2.23 2.46

0.074 0.168 0.434 0.859 1.46 0.251 0.789 1.09 1.42 1.70 0.475 1.06 1.67 1.97 2.32

4.49 4.37 4.18 4.01 3.84 4.41 4.18 4.10 4.02 3.93 4.39 4.24 4.09 4.02 3.96

0.032 0.037 0.102 0.274 0.500 0.119 0.239 0.378 0.511 0.782 0.250 0.528 0.671 0.815 0.895

0.009 0.027 0.078 0.232 0.445 0.093 0.215 0.348 0.475 0.735 0.200 0.468 0.812 0.773 0.875

4.71 4.39 4.17 3.91 3.72 4.32 4.16 4.02 3.92 3.71 4.29 4.08 4.00 3.90 3.80

0.057 0.169 0.499 0.893 1.31 0.254 0.917 1.23

0.034 0.118 0.378 0.758 1.14 0.201 0.797 1.11

4.71 4.48 4.22 3.99 3.80 4.58 4.28 4.16

0 0 0 0 1.91 0 0

W

5.36 5.82 5.74 5.61 4.85 5.23 4.85 5.23 4.76 4.19 5.71 4.76 4.57 4.76 3.33

i3QB

14.3 13.9 12.9 11.5 9.56 13.4 13.4 12.4

Solubility of AI in soils containing

organic

4175

matter

Table 6. Continued.

[Al3+]

[H+l

Added [OH-]

-Final[Na+] mM

Run #l (continued) 0.956 1.91 0 6.69 2.87 0.956 0 4.78 1.88 3.75 10.3 0 1.88 1.88 6.56 0 1.88 0 0 4.69 1.88 0.938 0 2.81 1.88 1.86 0 0.938

U’JOs-1

11.5 10.5 12.2 10.3 10.3 9.38 8.44

rAltotla [&lb

PH

1.54 1.78 0.523 1.20 1.83 2.13 2.36

1.41 1.67 0.429 1.04 1.69 2.01 2.33

4.05 3.93 4.56 4.37 4.20 4.08 3.98

NDC ND ND ND ND ND ND

0.078 0.162 0.435 0.838 1.39 0.063 0.157

4.51 4.36 4.19 4.01 3.83 4.51

Run#2 0 0 0 0 0 0 0 0

14.2 13.2 11.3 8.51 4.73 9.64 8.67 6.75

SkVBS 14.2 13.7 12.8 11.3 9.46 9.64 9.16 8.19

0 0

0 0.473 1.42 2.84 4.73 0 0.482 1.45

0

2.89

0

3.86

6.75

0

4.82

0

0

4.82

ND

0.435

ND ND

0.849 1.44

4.36 4.15 3.98

3.82

aConcentrationdetermined by AAS. bconcentrationdetermined according to CLARKE etaL(1992). CND = not determined.

agreed well with measured values. A regression analysis showed a highly significant linear relationship (r2 = 0.88, n = 8) between the calculated (A&+,) and measured increases in the Si concentration (ASi,,,) according to AS&,,, = 16( 2142) + 1.09( 50.40) ASi,,,,,

(7)

where concentrations are in yM, and the values in parentheses represent the 95% confidence interval. Obviously, the increase in reactive Al exceeds the estimated release of Al from silicate weathering (Table 5). This suggests that among the slowly reacting mineral phases the secondary non-silicatebound Al fractions dissolved fastest. 4.2.4. Experiment 4: Acid-base batch titrations in the presence of added aluminum

A final experiment was conducted to investigate how the increase in reactive Al (Experiment 3) influenced the relationship between H’ and Al’+ activities in solutions. Batch titrations were performed as in Experiment 2, but with additions of small aliquots (0.2 or 0.4 mL) of 0.025 or 0.05 M Al(N07)s solutions (reaction time 5 days). The Al( NO3)3 solutions were added as a final step after the NaNO, and HNOX/NaOH additions. Samples were shaken manually immediately after the Al(N03), additions to ensure a rapid dilution of the Al(NO,), solutions. Aluminum was added in amounts similar to the increase in CuCI, extractable Al between days 5 and 3 1 (Table 5), as well as twice this amount. For each soil horizon the ionic strength was held reasonably constant in all preparations. For the SkV Bh and Ros A soils, the ionic strength obtained was the same as that in Experiments l-3 (ca. 0.006 M), whereas for the SkV Bs and Ros

B soils, the ionic strength used in Experiment 4 was slightly higher (ca. 0.014 M). Experimental conditions are given in detail in Table 6. As shown in Fig. 3, an increase in the amount of reactive Al in the suspensions caused a shift in the pH-pA1 curves of all the samples in accordance with Eqn. 6; i.e., at higher levels of reactive Al the curves shifted to the right. What is particularly important, however, is that the observed shift in the pH-pA1 relationship for the SkV Bh and Ros A samples, which occurred between 5 and 30 days of equilibration, could be mimicked successfully by adding 5 pmol of Al (Fig. 3a,c) ; i.e. the amount of Al equivalent to the observed increase in CuC&-extractable Al between days 5 and 3 1 (Table 5). This result, which indicates that the Al’+ activity at equilibrium depends on the amount of reactive Al, would be unlikely in systems where mineral phases (e.g., AI(O control the Al solubility. Furthermore, the result supports our hypothesis that Al in solution is controlled by equilibrium reactions involving SOM in these soils. The pH-pA1 relationship for SkV Bs and Ros B after 5 days of equilibration in Experiment 4 does not compare well with that observed in Experiments 1 and 2, because of storage effects in Experiment 4 (see below). Therefore, only data from Experiment 4 are reported in Fig. 3b,d. 4.3. Simulations with WHAM The chemical equilibrium model WHAM (Tipping and Hurley, 1992; Tipping, 1994) was calibrated, using observations made in Experiment 4 with the SkV Bh soil. The model describes protonation and metal complexation reactions on discrete sites of humic substances. Recently, WHAM was successfully tested on organic soil horizons (Tipping et al.,

D. Berggren and J. Mulder

4176

5.0

(W

4.6

g

4.0

3.6

3.2 * 3.0

3.5

4.0

3.0 3.6

4.5

3.8

4.0

4.2

4.4

4.6

4.8

PH

PH 5.5

(cl

(d)

5.0 4.5 5 4.0 3.5

3.6

3.8

4.0

4.2

4.4

4.6

4.8

4.0

4.6

6.0

PH

PH

FIG.3. pA1 vs. pH for suspensions of (a) SW Bh, (b) SW Bs, (c) Ros A, and (d) Ros B samples. Data illustrate the effect of an addition of 5, 10, or 20 pm01 Al per 2.0 g field-moist soil at a constant equilibration time (5 d). Also illustrated for the SkV Bh and Ros A samples, is the effect of reaction time for systems with no added Al. For the SkV Bh and Ros A samples, data are from Experiments 1, 2, and 4. For the SW Bs and Ros B samples, only data from Experiment 4 are-presented owing to storage effects (see text). For Experiment 4 the chemical composition of the added solutions as well as pH and concentrations of A1,, and AL,, in the equilibrated suspensions are given in Table 6. The solid line indicates the solubility of gibbsite (log *KS0 = 8.85 at 8°C).

Table 7. Selected constants and parameters in the chemical equilibrium model WHAM describing the H+ and Al binding by humic acids (HA) and fulvic acids (FA). Values according to the data set SSED version 1.O (TIPPING,1994).

Constants/parameters mol type A groups per kg (dissociating at pH < 7) mol type B groups per kg m (dissociating at pH > 7) central pKfor type A groups PKA range of type A pK’s &KA central pKfor type 6 groups PKB range of type B plcs APKE ._ ._ PKA~HAcomplex formation constant tor type A groups P defines a charge interaction factor operating on pKs fraction of sites in bidentate f, configurations

HA

FA

3.29

4.73

1.65 4.02 1.76 6.55 3.43

2.37 3.26 3.34 9.64 5.52

1.3

0.4

-374

-103

0.5

0.4

nA

1995). Because our data suggest that organic complexation reactions also control the Al solubility in suspensions of our mineral soil horizons at pH < 4.1, we tested the WHAM model on one of them. Model calibration included optimization of the contents of humic acids (HA), fulvic acids (FA) , and reactive Al in the soil. The optimization procedure involved minimization of the difference between observed and calculated pH and Al,r, in solution. Constants and parameters used to describe the H + and Al binding by HA and FA are given in Table 7. For a complete description of WHAM, including the thermodynamic data for inorganic compounds, reference is made to Tipping ( 1994). In Fig. 4 we show the results of the calibration of WHAM for the SkV Bh titration data (Experiment 4). The optimized values for concentration of humic acids (CHA) and concentration of fulvic acids (CFA) (Table 2) suggest that the organic matter involved in complexation reactions constituted 49% of the Na&O, extractable organic matter and only 3 1% of the total organic matter. Similar fractions of organic matter, actively involved in metal complexation, were previously re-

4177

Solubility of Al in soils containing organic matter

(4

/ &

,,...i

, /r _,,;. ,”

3.6 : , i.....’ 0

,/,

3.2 : 3.0 -3.0

,

/

-2.0

-1.0

0.0

1.0

IbaseHacidl 1

0.8 . 0.6

-

??OAI A 5AI ??10 Al

0

9 .s -L

0.4 -

57

o2

’

1

’

,

(

( 2.0

: 3.0

(mM) ’

f

’

(

04

7 A \

“"..# '... \

“‘y

\

x.

curve for the SkV Bs and Ros B samples. This made a direct comparison of Experiment 4 with Experiments 1 and 2 difficult. Because the storage-induced change was most pronounced for SkV Bs, data for this horizon will be used to illustrate the storage effect. Storage shifted the pH-pA1 curve towards lower log Q values (Fig. 5 ) , with solutions approaching equilibrium with gibbsite at pH > 4.2. Adding a given amount of acid resulted in a lower pH and a higher pA1 for the stored sample (Experiment 4) than for the “fresh” sample (Experiment 2). “Fresh” and stored samples with the same acid addition and ionic strength are connected by arrows. It should be stressed that the samples had been carefully homogenized before they were frozen, and, as illustrated in Fig. 5 (Experiment 4: I = 14 mM), the reproducibility of the experimental procedure was good. Also illustrated in Fig. 5 is the minor effect of ionic strength on the pH-pA1 relationship in the range 0.010-0.014 M. The nature of the alteration of the SkV Bs and Ros B samples during the 11 ( 14) weeks of storage at 6°C is obscure. The fact that the SkV Bs and Ros B samples, but not the SkV Bh and Ros A samples, had a reduced Al solubility due to storage may indicate that the change was associated with structural changes in the reactive Al (OHb component.

'.. ?? o.o-3.0

*

'

-2.0

'

'

I.

'

-1.0

0.0

IbaseHacidl

1.0

5. DISCUSSION

I

2.0

3.0

(mM>

FIG. 4. Calibration of the mechanistic model WHAM for the SkV Bh titration data (Experiment 4). Symbols represent the experimental data. The lines are fits to WHAM-generated data; the solid line corresponds to zero addition of Al, the broken line to a 5 firno1 Al addition, and the dotted line to a 10 pmol Al addition. [AlJ is the

In Fig. 6 we present a conceptual model which may explain how the solubility of Al was controlled in our experiments. Aluminum is released simultaneously from rapidly reacting solid phases, quickly reaching an equilibrium state, and slowly reacting solid phases, not reaching equilibrium. Double arrows in Fig. 6 indicate equilibrium reactions and single arrows kinetically constrained nonequilibrium reactions. Chtr data suggest the presence of two kinds of slowly reacting solid

concentration of quickly reacting Al. The chemical composition of the added solutions as well as pH aad concentrations of Al,,,,and Ah, in the equilibrated suspensions are given in Table 6. Parameters and constants used in the simulations

0 /=lOmM(Exp.2) 0 1=10mM(Exp.4) II I = 14 mM (Exp. 4)

are given in Table 7.

ported for organic soil layers (Tipping et al., 1995). The optimized amount of reactive Al (CAl) for the SkV Bh sample exceeds the measured amount of exchangeable Al and is 71% of the amount of CuCIZ-extractable Al. The calibration curve for pH and AI,, against added base minus acid (Fig. 4) indicates reasonable fits for both observations. This successful calibration of WHAM on the SkV Bh data lends further support to the hypothesis that Al-organic matter interactions determine the Al solubility in strongly acidic mineral soil horizons,

1.5

A

4

??

0

3.0 $

Gibbsite

5.0 &_ 0 0

3.6

3.8

4.0

4.2

4.4

4.6

4.8

PH 4.4. Effect of Sample Storage In Experiments 1 and 2, samples were used immediately after thawing, whereas prior to Experiment 4, they were refrigerated for 11 or 14 weeks after thawing. For the SkV Bh and Ros A samples, in which the Al solubility was probably controlled solely by interactions with SOM, the 11 (14) weeks of storage at 6°C did not affect the pH-pA1 relationship obtained, and all three experiments resulted in about the same curve. By contrast, storage did affect the pH-pA1 equilibrium

FIG. 5. Effect of storage on the relationship between pH and pA1 for the SkV Bs soil. In Experiment 2, samples were used immediately after thawing, whereas in Experiment 4 samples were stored at about 6°C for 11 or 14 weeks after thawing. Numbers represent the concentration of H+ (in n&l) in the added solutions, and arrows connect data points obtained under the same experimental conditions for stored and “fresh” samples. The data points for Experiment 4, I = 14 mM, represent experiments performed on two different occasions, 11 (Run #1 in Table 6) and 14 weeks (Run #2 in Table 6) after thawing. The solid line indicates the solubility of synthetic gibbsite (log *I&, = 8.85 at 8°C).

4178

D. Berggren and J. Mulder

SLOWLY REACTING SOLID PHASES

I

J AP+( aq)

/\ RAPIDLY REACTING SOLID PHASES

w1

I

piijgcq

FIG. 6. A conceptual model describing the various pools of Al contributing to the Al solubility dynamics in the batch experiments. Double arrows indicate equilibrium reactions and single arrows kinetically constrained nonequilibrium reactions.

phases, i.e., primary and secondary Al-containing minerals. For the SkV Bh and Ros A samples, Al bound by SOM seems to be the only rapidly reacting solid phase present controlling the Al solubility. Not only were the suspensions of these samples strongly undersaturated with respect to any known Al (OH), phase, but also the H + -Al ‘+ relationship was far less than cubic. The latter also indicates that simple cation exchange did not control the solubility of Al. Support for the importance of SOM in regulating the Al”+ activity was also obtained in the Al addition experiment (Experiment 4)) where the Al solubility was shown to depend on the amount of reactive Al in the system. Further support for the importance of AI-SOM interactions in controlling soluble Al was obtained from titrations in the presence of different amounts of reactive Al, which could be reasonably well described with the mechanistic model WHAM (Tipping, 1994). In addition to a large pool of organically bound Al, the SkV Bs and Ros B samples seemed to include some form of rapidly reacting A1(OH)3, having a solubility which was slightly higher than that of gibbsite (log *KS, = 8.85 at 8°C; Palmer and Wesolowski, 1992). Observed log Q values were similar to those previously encountered in laboratory experiments with Spodosol Bs horizons conducted at 10°C (Dahlgren et al., 1989). For the SkV Bs and Ros B samples reducing pH below 4.1 resulted in a depletion of the Al( OH)? pool. In such conditions the activity of Al’+ was controlled by complexation reactions with SOM. For the Spodosol Bs horizons from the Hubbard Brook and Bear Brook sites Dahlgren et al. (1989) and Dahlgren and Walker (1993) proposed that the solution activity of Al ‘+, was controlled by hydroxy-Al interlayers of smectite. However, this does not seem to hold for the SkV Bs and Ros B horizons. The clay fraction in the SkV Bs and Ros B horizons was dominated by hydroxy-interlayered vermiculite (Table 3), and no detectable amounts of smectite were found. In hydroxy-interlayered vermiculite, Al seems to be tightly bound in the interlayers as highly positively charged polymers rather than as Al( OH), (Hsu and Bates, 1964). Therefore, we propose that surface precipitates of amorphous Al( OH)1 were the Al-controlling solid phase in the SkV Bs and Ros B samples at pH > 4.1. The *KS, values obtained in our study were about two orders of magnitude lower than those for amorphous Al(OH)? (e.g.,

Schecher and Driscoll, 1987). This may have been due to impurities of coprecipitated organic acids. Tipping et al. ( 1988) showed that the *KS,, of amorphous Al( OHh is lowered substantially in the presence of humic substances, the effect being stronger at higher concentrations of humic substances. 6. CONCLUSIONS This study suggests that the solubility of Al in many acidic mineral soil horizons, which are undersaturated with respect to macrocrystalline gibbsite, are controlled by complexation reactions with soil organic matter. This conclusion is based on results from batch experiments with a Spodosol Bh and an Inceptisol A horizon, which showed that (1) the H+-A17+ relationship in equilibrium solutions was far less than cubic, (2) the solubility of Al was strongly dependent on the amount of reactive Al in the system, and (3) the titration data in the presence of different amounts of reactive Al could be reasonably well described by the mechanistic model WHAM (Tipping, 1994). For a Spodosol Bs and an Inceptisol B horizon, batch titrations indicate the presence of a reactive Al (OH), phase regulating the activity of Al?+ in solutions at pH > 4.1, i.e., at pH values normally found in the field in these soil horizons. The solubility of this AI( phase appears to be slightly higher than that of gibbsite (log * KS0 = 8.85 at 8°C; Palmer and Wesolowski, 1992). However, this reactive Al phase may be present in small amounts and thus become easily depleted upon addition of large amounts of strong acid. At pH < 4.1, complexation reactions with soil organic matter seemed to control the Al” activity in these horizons. Our results illustrate the importance of considering the release of Al from slowly reacting mineral phases when interpreting Al solubility data obtained from batch experiments where soil samples are equilibrated for longer periods in acid solutions. This is particularly important when Al solubility is controlled by interactions with soil organic matter, because here the H + -Al’+ relationship is a function of the total amount of reactive Al in the system. Acknowledgments-Part of this work was performed at the Department of Ecology, Lund University, Sweden. We are grateful to S.

Solubility

of Al in soils containing

Billberg and E. Sjostrom, at Lund University, and to E. Velthorst and N. Nakken, at Wageningen Agricultural University, for help with the chemical analyses. X-ray diffraction analyses were performed by J. van Doesburg, Wageningen Agricultural University. We thank Ed Tipping for assistance with the WHAM applications. DB held a postdoctoral fellowship at the Swedish Council for Forestry and Agricultural Research and received a research grant from the Swedish Environmental Protection Agency. JM received a fellowship from the Royal Netherlands Academy of Arts and Sciences and had a research grant from the Dutch Additional Programme on Acidification, Phase III (APV III). D. J. Wesolowski and two anonymous reviewers are gratefully acknowledged for their comments on a previous version of this manuscript.

Editorial handling: T. Paces

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