iron speciation of in-situ solutions from an episaturated spodosol
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This article was downloaded by: [University of California Santa Cruz]On: 31 October 2014, At: 20:31Publisher: Taylor & FrancisInforma Ltd Registered in England and Wales Registered Number: 1072954 Registered office: MortimerHouse, 37-41 Mortimer Street, London W1T 3JH, UK
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Iron Speciation of In-Situ Solutions from anEpisaturated SpodosolElizabeth A. Rochette a & R. J. Cleary ba Washington State Department of Ecology , Kennewick, Washington, USAb Department of Natural Resources , University of New Hampshire , Durham, NewHampshire, USAPublished online: 05 Feb 2007.
To cite this article: Elizabeth A. Rochette & R. J. Cleary (2004) Iron Speciation of In-Situ Solutions from an EpisaturatedSpodosol, Communications in Soil Science and Plant Analysis, 35:3-4, 345-367, DOI: 10.1081/CSS-120029717
To link to this article: http://dx.doi.org/10.1081/CSS-120029717
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COMMUNICATIONS IN SOIL SCIENCE AND PLANT ANALYSIS
Vol. 35, Nos. 3 & 4, pp. 345–367, 2004
Iron Speciation of In-Situ Solutions from an
Episaturated Spodosol
Elizabeth A. Rochette1,* and R. J. Cleary2
1Washington State Department of Ecology, Kennewick,
Washington, USA2Department of Natural Resources, University of New Hampshire,
Durham, New Hampshire, USA
ABSTRACT
Complexation of iron (Fe) by organic compounds as well as reductive
dissolution are potentially important Fe dissolution processes in
organic-rich soils such as Aquods. The objectives of this study were to
develop a method to sample Aquod water table solutions and
determine solution Eh and pH values without introducing oxygen
from the atmosphere, and to determine the quantities of reduced and
organically-complexed Fe in the soil solutions. Two soil profiles of
a Typic Epiaquod in New Hampshire were studied. Zero tension
lysimeters were installed in each of 3 horizons of the profiles. Iron (II)
and total Fe were determined with the FerroZine method. Complexed
*Correspondence: Elizabeth A. Rochette, Washington State Department of
Ecology, 1315 W. 4th Ave., Kennewick, WA 99336-6018, USA; E-mail:
345
DOI: 10.1081/CSS-120029717 0010-3624 (Print); 1532-2416 (Online)
Copyright & 2004 by Marcel Dekker, Inc. www.dekker.com
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(including organically-complexed) Fe was assessed using anion
exchange and C18 columns. The solution samples were maintained
in a closed system throughout sampling and determination of Eh and
pH. In general, the redox potentials in the soil decreased with time
during the spring, while the pH values increased with time. Samples
obtained just before the Bs horizon drained had high levels of total
Fe(II) andC18-extractable Fe.Median percents of anionic Fe were 43,
12, and 0% in OA, Bhs, and Bs horizon solutions, respectively, while
those of C18-extracted Fe were 28, 66, and 0%. C18-extracted Fe,
total Fe and Fe(II) were all strongly related with linear r2 ranging from
0.92 to 1.00. This study indicates that Fe reduction can occur in spodic
horizons, and organically-complexed Fe(II) can be a significant form
of Fe in the soil solutions prior to draining of the profile. The latter
form of Fe should be considered in predicting Fe dissolution reactions
in Aquods.
INTRODUCTION
Iron cycling in soils can substantially influence trace element andphosphorus mobility, sulfur cycling, plant nutrition, soil morphology,and the practice of wetland delineation. Iron can precipitate with traceelements, sulfur (S) and phosphorus (P), and its oxides and hydroxidescan sorb these elements. Consequently, conditions that lead to dissolu-tion of Fe minerals can release associated elements to soil solutions.Furthermore, soil classification is based to some extent on extractablelevels of Fe in soil and soil colors influenced by Fe minerals. Though Fechemistry in soils has been extensively studied, the complex cycle of Fewith regard to the types and roles of organic compounds has not beenfully described.
Three main processes may lead to dissolution of Fe minerals in soils:reduction, organic chelation/complexation, and acid dissolution.[1] FerricFe oxide/hydroxide minerals are generally less stable under reducingconditions than under oxidizing conditions, and are thus prone toreductive dissolution. Soils that experience periodic flooding can becomereducing as a result of microbial consumption of carbon and oxygen. Theredox potential at which Fe reductive dissolution occurs depends on thepH, the form of ferric Fe present, and the types and amounts of solutes inthe soil solution. Low pH conditions can destabilize ferric Fe oxides andhydroxides even in oxidizing soils. Additionally, microbial activity cancatalyze reduction of ferric Fe minerals. Several Fe reducing microorgan-isms, such as Clostridium pasteurianum,[2] Shewanella putrefaciens,[3] andShewanella algae[4] have been observed to catalyze ferric oxide/hydroxide
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reduction. Consequently, predicting the stability of Fe in soil solutions iscomplicated by several environmental factors.
Complexation of Fe by organic compounds should be considered insoils that have abundant organic matter, such as Spodosols. Eluviation ofmetals such as Fe and aluminum (Al) from surface organic horizons to Bhorizons is generally thought to occur by either transport of metal–organiccomplexes through the soil profile, or mineral weathering and transport ofAl and silica as inorganic phases.[5] The processes leading to accumulationof organic matter in Bh horizons are not fully understood. Sorption and/orprecipitation of organic metal complexes due to decreasing carbon tometal ratios in metal-organic complexes, and/or microbial breakdownof organic components of the complexes followed by precipitation ofinorganic forms of Fe and Al have been proposed.[5] The nature ofthe organic compounds is not fully resolved, though carbohydrate andphenolic compounds have been shown to be important components ofdissolved organic carbon in soil solutions from Spodosols.[6] Low-molecular-weight organic acids typically account for less than 30% ofdissolved organic carbon, but can complex approximately 8–60% ofFe in solutions from Spodosols.[7–9] Phenolic compounds such asprotocatechuic, p-hydroxybenzoic, vanillic, and cinnamic acid have beensignificant in extracts from Spodosol samples,[10,11] and are thought toplay a role in Fe translocation in soils. The remainder of water-solubleorganic compounds is often reasoned to be fulvic and humic acids.
Aquods are Spodosols characterized by aquic conditions within50 cm of the soil surface, and a histic (seasonally saturated and organic)surface horizon or albic or spodic horizons with redoximorphicfeatures.[12] Thus, Fe cycling in Aquods is likely to be influenced byboth chelation and redox processes. In spodic horizons, podzolizationand reduction can be competing processes, with podzolization depositingFe and reduction dissolving Fe. Also, spodic horizons must pass thecriterion of organic matter illuviation as indicated by the optical densityof an oxalate extraction.[12] Some Aquods appear to be produced byhydrologic processes and are consequently low in Fe, but contain distinctBh horizons,[13] while others can have distinct zones of Fe accumulationin Bs horizons with less evidence of a hydrologic influence.[14] Theseasonally high water table may serve to limit downward translocation ofmetals relative to that in well-drained Spodosols.[14] Aquods provide anopportunity to examine the role of organic compounds and redoxprocesses in the cycle of Fe in soils.
The solution phase of Aquods could provide clues about redox andorganic complexation processes. Speciation of solution-phase Fe in fieldsoils can be partially addressed by the use of solid phase extraction.
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For Fe speciation, an assumption is that ‘‘free’’ Fe2þ and Fe3þ in solutionwill pass through anion exchange and hydrophobic matrices. Fe thatsorbs to anion exchange and hydrophobic matrices is not free Fe2þ andFe3þ, but complexed with ligands (organic or inorganic) in solution.Solid phase extraction approaches can provide a first approximation ofthe amount of Fe that is organically complexed in soil solutions. Theinitial hypothesis of this study was that organically-complexed Fe wouldbe the dominant form of Fe in solutions from organic-rich horizons, andthat reduced Fe would be the dominant form of Fe in solutions from thelower portion of the soil profile. One of the objectives of this study was todevelop a method to sample Aquod water table solutions and determinesolution Eh and pH values without introducing oxygen from theatmosphere. A second objective was to use the FerroZine method andsolid phase extraction to determine the valence and extent of organiccomplexation of Fe in the solutions. In addition, Fe and Al mineralequilibria are discussed.
MATERIALS AND METHODS
Site Description
Two soil profiles of the Squamscott Series, which is a sandy overloamy, mixed, active, mesic Typic Eqiaquod,[15] were studied. Theprofiles were located approximately 0.2 km west of the Squamscott River,near Newfields in Rockingham County, New Hampshire. The locationswere selected for their thick Bhs horizons, which would easilyaccommodate lysimeters. The horizons had the following thickness: Oe(4 cm), Oa/A (8 cm), E (23 cm Profile 1, 19 cm Profile 2), Bhs (14 cmProfile 1, 20 cm Profile 2), Bs (18 cm Profile 1, 16 cm Profile 2). The upperboundary of the C horizons was 61 cm in both profiles.
Soil Characteristics
Approximately 5 kg of each soil horizon were collected from pit facesprior to lysimeter installation. Samples were placed in plastic freezer bagsand frozen until use. Upon thawing, the samples were sieved moistthrough a plastic 2-mm sieve. Moist samples were used for extraction,and results were corrected for soil moisture (determined after ovendrying, 105�C for 24 h). Soils from the profiles were characterized with aselective sequential extraction procedure based on that of Shuman.[16]
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However, the bleach extraction step was replaced with a sodiumpyrophosphate extraction procedure described by Loeppert andInskeep.[17] One gram of each sample was extracted with (1) 15mL of1M Mg(NO3)2 (pH 4.6, consistent with the pH of the soils), followed by(2) 25mL of sodium pyrophosphate, followed by (3) 50mL of 0.2Mammonium oxalate/oxalic acid (pH 3) carried out in the dark, followedby (4) 50mL of 0.2M ammonium oxalate/oxalic acid (pH 3) plus 0.1Mascorbic acid. The extractants were selected to extract the followingfractions of Fe, Al, and manganese (Mn) in the soil: (1) exchangeable,(2) organic-bound, (3) amorphous, and (4) crystalline. Extractions wereperformed in triplicate. Selective extractions such as this are known to beless selective than desired, but provide some indication of soil–metalassociations. Soil pHs were determined with a 1:1 deionized water:soilextraction. Characterization data are given in Table 1.
Lysimeters and Solution Collection
Lysimeters were prepared from 1.7-L polypropylene containers,26 cm� 15.2 cm� 3.8 cm. The containers were trimmed to taper alongtheir length to a height of 1.3 cm, so that the ‘‘back’’ edge was shallowerthan the ‘‘front.’’ Two holes were drilled in the front side of the lysimetersto accommodate quarter-inch (0.64 cm) components of LPDE quick-connects. Nylon screening was placed over the inlet sides of the quick-connects to keep large soil particles from flowing out of the lysimetersinto the tubing; the screening was fixed in place with plastic zip-ties.Quarter-inch (i.d.) reinforced PVC tubing was attached to the outlet endsof the quick-connects (two sections, 20 cm long). The two portions werejoined with a quarter-inch Y connector. Additional tubing (28 cm) wasattached to the Y connector to provide a single outlet for each lysimeterto the collection devices. All fittings and tubing for the lysimeters wereobtained from VWR Scientific (VWR Scientific, Bridgeport, NJ).
To install the lysimeters, sampling pits (approximately 1m2� 0.7m)
were excavated and soil from each horizon was set aside on tarps. Threelysimeters were installed in each of the A, Bhs, and Bs horizons. Becausethe A horizon was thin, the lysimeters were installed at the base of the A(12 cm depth). The Bhs lysimeters were installed at 40 cm depth (Profile 1)and 41 cm depth (Profile 2), and the Bs lysimeters were installed at 57 cmdepth (Profile 1) and 58 cm depth (Profile 2). Different faces of theexcavated pits were used for installing the A, Bhs, and Bs horizonlysimeters, to prevent shallow lysimeters from interfering with deeperlysimeters. Blocks of soil approximately 38 cm thick and as deep as the
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bottom of the horizon of interest were removed from each face prior tolysimeter installation. Soil was then gently excavated using plastic knivesto allow lysimeters to be put in place flush with the soil above them,minimizing ‘‘dead space.’’ The lysimeters sloped such that the back ofthe lysimeters was approximately 1.3 cm higher than the front of thelysimeters, to facilitate water flow toward the outlets of the lysimeters.
Table 1. Soil characterization data for soils from Profile 1 and 2. Standard
deviations are given in parentheses.
Horizon/profile OA/1 E/1 Bhs/1 Bs/1 OA/2 E/2 Bh/2 Bs/2
Moist color 10YR
2/1
2.5Y
4/1
2.5YR
2.5/1
5YR
3/4
10YR
2/1
2.5Y
4/1
10YR
2/1
7.5YR
3/4pH 3.80 4.26 4.32 4.39 3.75 4.00 4.25 4.39
Fe (mmol kg�1)
Mg(NO3)2 0.16
(0.00)
0.02
(0.00)
0.05
(0.01)
0.02
(0.00)
0.06
(0.00)
0.03
(0.01)
0.06
(0.03)
0.05
(0.01)Na pyrophosphate 30
(1)
1.0
(0.1)
9.9
(1.0)
62
(11)
14
(1)
1.8
(0.1)
7.9
(0.9)
12
(1)AODa 3.5
(0.2)
0.18
(0.01)
7.9
(1.1)
135
(22)
0.93
(0.49)
0.22
(0.05)
6.0
(1.6)
5.8
(1.7)AOþAscorbic acid 7.0
(0.7)
1.6
(0.2)
28
(1)
295
(16)
1.9
(0.4)
0.75
(0.06)
29
(1)
23
(0)Al (mmol kg�1)
Mg(NO3)2 13
(0)
4.2
(0.0)
3.9
(0.1)
1.3
(0.1)
13
(0)
5.4
(0.2)
10
(0)
3.7
(0.6)Na pyrophosphate 48
(3)
7.6
(0.2)
89
(10)
58
(4)
32
(2)
6.6
(1.5)
198
(7)
94
(3)AODa 12
(0)
1.8
(0.1)
20
(2)
39
(4)
3.5
(0.1)
1.3
(0.1)
15
(2)
85
(2)AOþAscorbic acid 7.6
(0.5)
2.6
(0.4)
21
(1)
52
(4)
2.6
(0.1)
1.3
(0.0)
27
(0)
20
(0)Mn (mmol kg�1)
Mg(NO3)2 46
(4)
3
(1)
4
(1)
4
(0)
9
(0)
1
(0)
2
(0)
2
(2)Na pyrophosphate 31
(0)
2
(1)
24
(2)
431
(72)
9
(3)
2
(0)
11
(0)
15
(5)AODa 15
(0)
NDb 101
(26)
755
(118)
2
(2)
ND 19
(5)
31
(2)AOþAscorbic acid 35
(2)
10
(2)
170
(5)
795
(46)
9
(4)
3
(1)
171
(1)
137
(30)
aAmmonium oxalate extraction performed in the dark.bBelow detection level.
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Soil from the horizon excavated was packed beneath each lysimeter tomaximize soil-lysimeter contact. The soil above the pan portion ofthe lysimeters was undisturbed except at its immediate contact with thelysimeters. The tubing of the lysimeters was buried using soil from thehorizon of interest. A wall of soil approximately 38 cm thick (measuredfrom the pit face toward the lysimeter) was then repacked using eachhorizon to its original depth above the tubing of each lysimeter to rebuildthe soil profile above the tubing. The walls of repacked soil were held inplace with pegboard and wooden stakes. Approximately 10 cm of tubingwere allowed to extrude from the pegboard for solution collection. Thetubing was filled with deionized water and capped during the dry season.During wet seasons, the tubing was filled with soil water and capped.
Samples were collected when the water table was above the lysimeters.Flow was initiated by pumping water out of the sampling pits to a depthjust below the outlets of the tubing in the horizon sampled. Water waspumped using a sump pump powered by a generator. Water removed bypumping was routed with a hose to a location roughly 20m away fromthe sampling pits. Approximately 20mL of water was allowed to flow outof the lysimeter tubing prior to connection of the sampling bags. Thelysimeter outlets were connected to 1-L Tedlar sampling bags (SKC,Eighty Four, PA) during collection. Air bubbles in the bag were removedwith a syringe immediately after collection. Approximately 0.3 L ofsolution was collected in each bag. The plastic valves on the bags werethen closed, and the samples were placed in a dark box for transport tothe laboratory. Tedlar bags are impermeable to gases; we have foundFe2þ solutions to be stable in the bags for at least a week. An alternativesampling approach, attaching the outlets of the lysimeters to the Tedlarbags via a peristaltic pump, was tested and found suitable, though slower.However, the latter approach avoids the need for a sampling pit.Temperatures were obtained for each horizon immediately after samplecollection, with a Hanna Instruments HI 9053 and 82-cm K-typethermocouple thermometer (Ben Meadows, Inc., Janesville, WI).
Solution pH and Redox Potential
Solution Ehs and pHs were measured with an Orion 710A meter. ACorning platinum combination redox probe (Ag/AgCl reference) and anOrion combination pH probe were each placed inside glass closed celldevices obtained from Cole Parmer, Inc. (Vernon Hills, IL). Sampleswere removed from the sampling bags with syringes and filtered throughin-line 0.45 mm filters during transfer to the cells. A new filter was used
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for each sample. Plastic valves were installed on inlets and outlets of theclosed cells, and were closed while measurements were taken. Excesssample was pushed through the cells as necessary to assure that therewere no air bubbles in the cells. The cells were rinsed copiously withdeionized water before and after introduction of each sample. Eh valueswere corrected to standard hydrogen electrode values at 25�C by adding199mV.[18] The Eh and pH values were laboratory measurements madefor comparison with laboratory speciation data from individual lysimetersamples, and are not interpreted as representing true field values.
Iron Analysis
All soil solutions were filtered with 0.45 mm filters to isolate Fe that istechnically ‘‘dissolved’’ and highly mobile in soils. Filtrates thereforeincluded some colloidal (>1000Da) Fe. Ultrafiltration was notperformed to further isolate ‘‘truly dissolved’’ Fe because charged orhydrophobic colloids as well as ‘‘truly dissolved’’ components were ofinterest. Ultrafiltration could precede FerroZine analysis and solid phaseextraction if size discrimination is desired. Iron (II) and total Fe weredetermined using the FerroZine method as described by To et al.[19] Iron(II) was determined in the absence of the reductant hydroxylaminehydrochloride, while total Fe was determined in the presence of hydroxyl-amine hydrochloride. To correct for interference resulting from coloreddissolved organic matter, samples were analyzed at the same wavelengthin the absence of FerroZine, and the absorbances were subtracted fromthose obtained in the presence of FerroZine (as described by To et al.[19]).
Solid Phase Extraction
The anion exchanger used in this study was Amberlite IRA-958(Sigma-Aldrich, St. Louis, MO), 16–50 mesh, macroporous, strong anionexchanger with quaternary ammonium functional groups on polystyrene.It was slurry packed in 1-cm diameter glass columns with 28 mm filter fritsfrom Bio-Rad (Hercules, CA). Prior to use, it was cleaned with 15mL ofthe following sequence of solutions, with a 10mL deionized water rinsebetween each solution: 0.5M NaOH, 0.5M HCl, 1M acetic acid, andacetate buffer of the desired pH (3–6). Columns were allowed toequilibrate with the acetate buffer over night, and were regenerated with15mL of acetate buffer between samples. Ten mL of filtered sample werepassed over the columns and collected in clean 20mL scintillation vials
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(gravity drainage). Ten mL of deionized water were then passed over thesame column and collected with the sample, in the same scintillation vial.Solutions were quantitatively transferred to 25mL volumetric flasks foranalysis by the FerroZine method for total Fe. Each column was used upto ten times before replacement.
Maxi-Clean C18 columns (strongly nonpolar, reverse-phase with50 mmparticles and 20 mm filter frit) were obtained fromAlltech Associates(Deerfield, IL). Each cartridge was activated by rinsing with 5mL ofmethanol, followed by 5mL of deionized water before samples wereintroduced. Five mL of filtered sample were passed over the columns andcollected in clean 20mL scintillation vials. Five mL of deionized waterwere then passed over the same column and collected with the sample, inthe same scintillation vial. Solutions were quantitatively transferred to25mL volumetric flasks for analysis by the FerroZine method for total Fe.A new cartridge was used for each sample, though it is probable thatcartridges could be regenerated with methanol and used again.
Aluminum and DOC
Filtered samples (0.45 mm) were acidified to pH 2 with HCl (OmniTrace, EM Science, Darmstadt, Germany) and analyzed for Al on a VistaAX/CCD Simultaneous ICP-AES by Varian (Palo Alto, CA). Sampleswere also analyzed for DOC, on a Shimadzu (Kyoto, Japan) 5000 TOCanalyzer. To prepare for DOC analysis, samples were filtered throughpre-ashed 0.7 mm Whatman (Maidstone, England) glass fiber filters intopre-ashed 40mL glass amber VOA vials. Samples were preserved byreducing the pH to approximately 2 with sulfuric acid.
RESULTS AND DISCUSSION
Water Levels and Redox Status of Solutions
As a result of a drought in the fall and winter of 2001–2002, waterlevels were below all lysimeters until late February, 2002. The water tablerose during March; the soil profile was essentially flooded on March 27,2002 (Tables 2 and 3). Water levels were dynamic throughout April andMay as a result of frequent rain events. Water levels dropped from May28 through July 2, after which water levels were below all lysimeters.Water levels in Profile 1 were generally higher than those in Profile 2(relative to the soil surface), thus, Profile 1 could be sampled over a
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longer period of time. Samples from the surface horizons (beneath the Ahorizon) were only available from March 27 through April 30, while theBhs lysimeters were sampled from March 12 through May 21. Thesampling period for Bs lysimeters was from March 8 through July 2.
The redox potential in all horizons generally decreased with timeduring the spring, while the pH values increased with time (Tables 2and 3). This is consistent with a progression toward reducing conditionsas the soil temperature increased during the spring. Rain eventsoccasionally interrupted this trend. For instance, a decrease in solutionpH values on April 17 (all horizons) was coincident with an increase inwater level on that date, which was brought about by heavy rainfall in theweek prior to sampling. On the other hand, the rapid rise in the watertable observed between March 8 and 12 appears to have resulted in adecrease in Fe in the Bs horizon solutions, along with an increase in bothEh and pH. This suggests dilution of the soil water by groundwater fromupslope sources, though direct snowmelt and rainfall inputs were alsolikely at this time. Solution pH values generally increase with depth, in
Table 2. Redox status and related characteristics of solutions from all sampling
dates (2002) and soil horizons in Profile 1. Values are means (standard deviations
in parentheses).
Date Hor.
Total Fe
(mmolL�1)
Fe(II)
(mmolL�1)
Redox
potential
(mV) pH range
Temp.a
(�C)
Water
levelb (cm)
3/27 OA 2.42 (0.25) 1.68 (0.07) 476 (15) 3.33–3.43 2.6 3
4/17 OA 2.56 (0.82) 1.83 (0.13) 464 (2) 3.50–3.70 10.6 4
5/7 OA 4.05 (0.04) 1.75 (0.50) 449 (6) 3.86–3.97 12.1 12
5/21 OA 3.97 (0.73) 2.85 (0.13) 465 (7) 3.93–3.98 10.3 6
3/12 Bhs 0.45 (0.09) 0.41 (0.09) 544c 4.34c 4.6 33
4/17 Bhs 1.40 (0.11) 0.84 (0.34) 473 (12) 3.71–3.80 8.2 4
5/7 Bhs 6.62 (5.07) 6.59 (5.91) 424 (0) 4.35–4.40 8.4 12
5/21 Bhs 12.9 (1.00) 12.9 (1.07) 415 (4) 4.45–4.54 9.3 6
3/8 Bs 0.57 (0.13) 0.43 (0.20) 466 (14) 3.78–3.90 4.0 42
3/12 Bs 0.18 (0.04) <0.18 563 (3) 4.43–4.49 5.3 33
4/17 Bs 0.38 (0.04) <0.18 495 (8) 3.91–4.04 6.9 4
5/7 Bs 1.09 (0.98) 0.66 (0.84) 416 (4) 4.59–4.69 7.7 12
5/28 Bs 0.54 (0.05) 0.36 (0.23) 417 (4) 4.28–4.30 9.9 20
7/2 Bs 49.1 (42.8) 44.8 (38.1) 411 (13) 4.36–4.87 13.9 50
aTemperatures represent the horizon/profile sampled.bWater level values were water depths beneath the surface in the profile sampled.cn¼ 1.
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the order OA<Bhs<Bs (Fig. 1). The seemingly high redox potentialsare low enough to allow some reduction of Fe at the low pH values ofthese soils; Fe oxides and hydroxides are more susceptible to reductionat low pH values. For instance, the equilibrium Eh for reduction ofamorphous Fe(OH)3 at pH 4.5 is as high as 614mV assuming a totaldissolved Fe activity of 10�6.[20]
The final samples, obtained on July 2, had high levels of Fe, largely Fe(II), suggesting that reductive dissolution of Fe minerals occurred in theFe-rich Bs horizon prior to its draining. However, reductive dissolutionof Fe minerals had somewhat variable impacts on the solution Feconcentration in the Bs horizon on this date, as total Fe concentrationsin solution ranged from 6.2 to 92 mmolL�1 for the three lysimeters sampledin Profile 1.
Solid Phase Extraction
Originally, DEAE-Sephadex, a weak anion exchanger, was tested asan anion exchanger based on the method of Camerlynk and Kiekens.[21]
Table 3. Redox status and related characteristics of solutions from all sampling
dates (2002) and soil horizons in Profile 2. Values are means (standard deviations
in parentheses).
Date Hor.
Total Fe
(mmolL�1)
Fe(II)
(mmolL�1)
Redox
potential
(mV) pH range
Temp.a
(�C)
Water
levelb (cm)
3/27 OA 1.68 (0.63) 1.25 (0.20) 484 (8) 3.37–3.72 4.6 8
4/2 OA 2.20 (0.77) 0.61 (0.20) 446 (2) 4.31–4.35 4.7 10
4/30 OA 1.16 (0.18) 0.86 (0.16) 447 (4) 3.74–3.88 7.4 7
4/9 Bh 0.45 (0.008) <0.18 459 (2) 4.35–4.45 6.6 18
4/23 Bh 0.64 (0.14) 0.57 (0.16) 454 (3) 3.95–4.07 8.9 20
4/30 Bh 0.79 (0.21) 0.70 (0.34) 455 (7) 4.05–4.24 7.7 7
3/8 Bs 0.47c 0.29c 491c 4.40c 4.0 45
4/9 Bs 0.27 (0.07) <0.18 467 (5) 4.47–4.67 6.5 18
4/23 Bs 0.27 (0.11) 0.21 (0.05) 462 (3) 4.27–4.36 8.9 20
4/30 Bs 0.36 (0.07) 0.29 (0.07) 468 (2) 4.42–4.53 8.0 7
6/4 Bs 0.73 (0.11) <0.18 420 (0) 4.62–4.67 11.5 51
aTemperatures represent the horizon/profile sampled.bWater level values were water depths beneath the surface in the profile sampled.cn¼ 1.
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However, it removed up to 67% of Fe3þ from ferric ammonium sulfatesolutions, and was abandoned. Instead, a strong anion exchanger,Amberlite (IRA-958), was found to sorb less than 40% of Fe in solution.The optimal pH of the regeneration buffer was found to be 5; this pHgave consistent recoveries nearing 85% for both Fe2þ and Fe3þ (Table 4).The C18 columns were also tested for Fe retention and were found toprovide 86(�1.6)% recovery of Fe(II), and 102(�3.5)% recovery of Fe3þ.The anion exchanger was then optimized as an anion exchanger usingphosphate. Removal of phosphate was improved from 92(�2.7) to100(�0.1)% by increasing the amount of resin used from 5mL to 7.5mL.
350
400
450
500
550
600
3 3.5 4 4.5 5
pH
Eh
(m
V)
Figure 1. Relationship between redox potentials and pH values for solution
samples from A horizons (^), Bhs horizons (g), and Bs horizons (i). Samples
from July 2 are included. (View this art in color at www.dekker.com.)
Table 4. Recoveries of iron (3.6–7.2 mmolL�1) from anion exchange columns.
Optimal resin volume was determined with a buffer pH of 5. Standard deviations
are given in parentheses.
Buffer pH
Fe species 3 4 5 6
Fe(II) 67 (5.3) 101 (17.5) 84 (4.6) 70 (3.8)
Fe(III) 53 (4.4) 77 (0.3) 86 (4.9) 67 (5.2)
Resin volume (mL)
Fe species 5 7.5
Fe(II) 84 (4.6) 93 (8.2)
Fe(III) 86 (4.9) 95 (3.9)
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Retention of Fe with the larger column volume was not increased by this
change. A subset of field solution samples from throughout the season
were spiked with Fe3þ to determine column efficiencies. Recovery of Fe3þ
from the anion exchange columns was 92(�25)% (median of 86%), and
that from the C18 columns was 94(�28)% (median of 87%). It therefore
appears that natural samples give more variable recoveries than do
spiked deionized water samples. Recoveries greater than 100% suggest
positive interferences from matrix components (such as metals) in the
natural samples, while values less than 100% suggest that some spiked Fe
was held by the columns, possibly as a result of interaction of the spiked
Fe with natural matrix components and the columns.Solid phase extraction results are presented graphically for repre-
sentative lysimeters from each horizon, each profile (Figs. 2–4).Anionic Fe generally increased with time in the OA and Bhs horizons
of Profile 1, and was near the detection level in the Bs horizon. In Profile 2,
however, it peaked on April 4 in the OA horizon, but was near detection
Date
Fe (µmol L-1)
0
1
2
3
4
5
3/25 4/14 5/4 5/24
a
0
1
2
3
4
5
3/25 4/14 5/4 5/24
b
Figure 2. Iron concentrations in solutions from representative lysimeters in the
OA horizons: lysimeter 2 in the Profile 1 (a) and lysimeter 2 in Profile 2 (b). The
symbols refer to total iron (g), Fe (II) (^), anionic (i), and C18-extractable (�).
(View this art in color at www.dekker.com.)
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level in the B horizons with the exception of June 4 in the Bs horizonwhen it reached its highest level. Anionic Fe was quite variable, especiallyin B horizon samples. The mean, standard deviation, and medianpercents of total Fe in the OA horizon samples were 37(�22) and 43%,respectively. In the Bhs horizon samples anionic Fe was 19(�19)% oftotal Fe (median of 12%). In the Bs horizons samples anionic Fe was26(�33)% of total Fe (median of 0%). It is possible in natural samplesthat some of the Fe retained by an anion exchange column is zwitterionic,having both a negative and a positive charge.
Because C18 is an organophyllic matrix, Fe extracted by C18columns was also assumed to represent organically-complexed Fe. Ironextracted with C18 in the OA samples paralleled that of total Fe, peakingon May 7 in Profile 1 and April 4 in Profile 2; it increased with time in Bhhorizon samples from both profiles; and it declined from March through
Date
Fe (µmol L-1)
0
5
10
15
3/5 3/25 4/14 5/4 5/24 6/13
0
0.4
0.8
1.2
3/5 3/25 4/14 5/4 5/24 6/13
a
b
Figure 3. Iron concentrations in solutions from representative lysimeters in the
Bhs horizons: lysimeter 2 in the Profile 1 (a) and lysimeter 2 in Profile 2 (b).
Concentrations of uncharged iron species were below detection levels (data not
shown). Concentrations below 0.18mmolL�1 for species are included for
completeness. Note that Profile 2 concentrations are graphed on a different
scale than those of Profile 1. The symbols refer to total iron (g), Fe(II) (^),
anionic (i), and C18-extractable (�). (View this art in color at www.dekker.com.)
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May in Bs horizon samples, and then increased after May as the watertable fell in both profiles. C18-extracted Fe was 24(�22)% of total Fe(median of 28%) in the OA samples, 52(�29)% (median of 66%) in the
Bh horizon samples, and 30(�39)% (median of 0%) in the Bs horizonsamples. Anionic plus C18-extracted Fe ranged from 57(�38)% (median
of 64%) of total Fe in OA samples, to 68(�35)% (median of 77%) in Bhshorizons, and 49(�56)% (median of 17%) in Bs horizon samples. Itappears that some Fe was present in surfactant-like organic complexes,
with charged regions and organophyllic regions, because some samplescontained greater than 100% anionic plus C18-extracted Fe. Icopiniand Long[22] observed that chromium(III) extracted by anionic and
hydrophobic columns exceeded 100% of available chromium in watersamples, and it was suggested that some Cr(III) may be present in forms
that are both anionic and hydrophobic.The concentrations of anionic and C18-extracted Fe were related to
each other as well as to those of total Fe and Fe(II) when data from all
Date
Fe (µmol L-1)
0
0.2
0.4
0.6
0.8
3/5 3/25 4/14 5/4 5/24 6/13
0
0.2
0.4
0.6
0.8
2/13 3/5 3/25 4/14 5/4 5/24 6/13
a
b
Figure 4. Iron concentrations in solutions from representative lysimeters in the
Bs horizons: lysimeter 2 in the Profile 1 (a) and lysimeter 1 in Profile 2 (b).
Concentrations below 0.18mmolL�1 for species are included for completeness.
Samples from July 2 are omitted. The symbols refer to total iron (g), Fe(II) (^),
anionic (i), and C18-extractable (�). (View this art in color at www.dekker.com.)
Iron Speciation of In-Situ Solutions 359
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horizons were combined. Anionic Fe generally related to C18-extractedFe, Fe (II), and total Fe only weakly (linear r2 from 0.38 to 0.47) whenJuly 2 was not included in the regression. Anionic Fe in the July 2samples was linearly related to C18-extracted Fe, Fe(II), and total Fe; ther2 values for regressions of anionic Fe against C18-extracted Fe, Fe(II),and total Fe for all samples ranged from 0.95 to 0.97 when July 2 wasincluded. C18-extracted Fe and Fe(II) were strongly related, with r2 of0.99 and 0.95 with and without July 2 samples, respectively. C18-extracted Fe and total Fe were also strongly related with r2 of 1.00 and0.92 with and without July 2 results, respectively (Fig. 5). Regressions ofthe sum of C18-extracted Fe plus anionic Fe against Fe(II) and againsttotal Fe yielded lower r2 values than did those of C18-extracted Fe,though no r2 was less than 0.84. Not surprisingly, total Fe and Fe (II)were closely related (linear r2¼ 0.99 and 0.95 with and without July 2included, respectively).
It is apparent from these results that organically-complexed Fe wasfrequently reduced Fe, and dissolved Fe was dominated by reduced Fe.This was especially true for samples with the highest levels of Fe.However, total, reduced, and C18-extracted Fe in samples from thecolder and hydrologically more active period of March through Aprilfrequently had separate temporal trends (Figs. 2–4). It also appears thatorganic complexation was quite important in the Bhs horizons and evenin Bs horizons, which is contrary to the original hypothesis. The Bhorizon samples were obtained over longer time periods and were morevariable than OA horizon samples. It appears that Fe in the B horizon
050
100150200250300350
0 50 100 150 200 250 300
C18-Extracted Fe (µmol L-1)
Total Fe (µmol L-1)
Figure 5. Relationship between C18-extracted iron and total iron. The
regression equation: Total Fe¼ 1.18 (C18-extracted Fe)þ 2.22, r2¼ 1.00. (View
this art in color at www.dekker.com.)
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solutions differs from that in the OA horizon solutions; the OA horizonsamples had greater relative proportions of anionic Fe, and lower relativeproportions of Fe(II) and C18-extracted Fe than did B horizon samples.Also, reduced Fe often exceeded 50% of total Fe in the OA solutions.
Aluminum and Dissolved Organic Carbon
The relationships between Al and DOC, and total Fe and DOC differfrom horizon to horizon (Figs. 6a and b). The OA horizon solutionsgenerally contained the highest concentrations of DOC (excluding July 2)but a relationship between Al and DOC did not exist, while therelationship between total Fe and DOC was weak (linear r2¼ 0.35) in thishorizon. The Bhs solutions had intermediate concentrations of DOCwhen compared to those of the OA and Bs horizons. The Bhs horizonwas the only horizon having notable relationships between Al and DOC(linear r2¼ 0.87) and total Fe and DOC (linear r2¼ 0.79). Sodium-pyrophosphate-extractable (organically-bound) Al was the dominantform of solid-phase Al in this horizon (Table 1), though crystalline formsof Fe were the dominant solid-phase forms in this horizon. Iron, and to alesser extent Al, were distinctly higher in Bhs solutions during May thanin earlier months. To allow closer examination of the Bs horizon samplesobtained from March 8 through June 4, samples from July 2 wereexcluded from regression analyses involving Fe, as July 2 samples had Feconcentrations up to 2 orders of magnitude greater than those fromearlier dates. In the Bs horizon, Al and Fe concentrations had essentiallyno relationship to DOC even though sodium-pyrophosphate extractableAl was the dominant form of solid-phase Al in this horizon. As the springprogressed, Fe concentrations in solution increased more dramaticallythan did those of Al and DOC as a result of reductive dissolution of Fe.C18-extractable Fe did not relate to DOC in any of the horizons.
Mineral Equilibria
The lowest redox potentials and highest Fe concentrations observedduring the sampling season were found in the samples from the Bshorizon, Profile 1, on July 2 (Table 3). Lysimeter 1 had a total Feconcentration 14 times lower than that of lysimeter 3 and approximately8 times lower than that of lysimeter 2. Most of the Fe on this date was
Iron Speciation of In-Situ Solutions 361
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Fe(II), while an average of 84% of the Fe was sorbed by C18 columns(Fig. 7). Roughly half of the Fe was sorbed by anion exchange columns.
Because of the high Fe levels, the samples from July 2 were selected formineral equilibria calculations and analysis of anions. It was hypothesizedthat the high Fe concentrations were indicative of reductive dissolution ofFe minerals. Ferrous Fe accounted for nearly 85 to 96% of the Fe in thesolutions on this date. Visual MINTEQ[23] was used to determine mineralsaturation indices, and species activities in solution. All three lysimetersolutions were supersaturated with respect to several common Fe oxideswith the exception of lysimeter 1, which was undersaturated with respect toferrihydrite and maghemite (Table 5). It is possible that ferrihydrite andmaghemite control the concentration of Fe in this soil.
The following explanations for supersaturation with respect to Fe(III)minerals are offered: (1) solid and solution phases, especially in lysimeters2 and 3, were not at equilibrium on this date (‘‘stable’’ solid phases were
0
10
20
30
40
50
60
70
0 1 2 3 4
02468
10121416
0 1 2 3 4
DOC (mmol L-1)
Total Fe (µmol L-1)
Al ( µmol L-1)
a
b
Figure 6. Relationships between aluminum and DOC (a) and total iron and
DOC (b) for solution samples from OA horizons (^), Bhs horizons (g), and Bs
horizons (i). Samples from July 2 are omitted.
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not precipitating as fast as reductive dissolution of Fe occurred); (2)
microbial activity catalyzed reductive dissolution under conditions not
predicted by equilibrium calculations; or (3) organic complexation limits
the extent to which ‘‘stable’’ minerals precipitate. Kinetically-limited
reductive dissolution of ferrihydrite has been described by Langner and
Inskeep.[24] However, kinetically-limited precipitation would be required
by explanation 1. Several studies indicate that Fe-reducing microorgan-
isms catalyze reduction of ferrihydrite.[2,4,24,25] However, it is not clear that
thermodynamics are violated as a result of microbial activity. McCreadie
et al.[27] inferred reductive dissolution of Fe oxide minerals (hematite and
maghemite) as a source of Fe(II) below the water table in mine tailings,
even when thermodynamic calculations indicated supersaturation with
respect to the Fe oxides. Fredrickson et al.[3] observed supersaturated
conditions for several minerals in a microbial/hydrous ferric oxide system.
In the same study, an organic buffer (PIPES) appeared to increase solution
concentrations of Fe(II) relative to solutions without PIPES, possibly
through complexation of Fe(II). In our study, adding DOC as MINTEQ
input (Gaussian model) had negligible influence on Fe mineral saturation
indices; however, without further examination of the specific organic
compounds present and consideration of Fe-organic stability constants for
the appropriate organics, explanation 3 cannot be ruled out.
TOTAL_FE
FE_II_
ANIONIC
C18_EXT0
20
40
60
80
100
Lysimeter 1 Lysimeter 2 Lysimeter 3
Fe µmol L-1
Figure 7. Iron concentrations in solutions from July 2, 2002 (Bs horizon, Profile 1).
Iron Speciation of In-Situ Solutions 363
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The concentrations of Al and DOC in samples from July 2 wereintermediate relative to those observed throughout the season. On July 2,Al concentrations ranged from 8.5 to 28 mmolL�1 (lysimeter 1> lysimeter2> lysimeter 3). Important Al minerals appear to be boehmite,basaluminite (Al4(OH)10SO4), amorphous Al(OH)3, jurbanite, andgibbsite. Lysimeter 1 (pH 4.36) was undersaturated with respect toboehmite, while lysimeters 2 (pH 4.65) and 3 (pH 4.87) were super-saturated with respect to boehmite. All lysimeters except lysimeter 3 wereundersaturated with respect to basaluminite, with saturation indicesranging from�2.119 to 0.102 (lysimeter 1< lysimeter 2< lysimeter 3).All lysimeters were undersaturated with respect to amorphousAl(OH)3, while supersaturated with respect to gibbsite and diaspore.DOC concentrations ranged from 200 to 233 mmolL�1 (lysimeter 3>2> lysimeter 1). Adding DOC as MINTEQ input (Gaussian model) hadminimal impact on mineral saturation indices with the exception of
Table 5. Visual MINTEQ output for Profile 1, Bs lysimeters on July 2. Solution
species with activities less than 10�8 are not given.
Lysimeter Solution species
Activity in
solution Mineral
Saturation
index
1 Cl� 1.58� 10�4 Ferrihydrite �0.29
Fe ðOHÞþ2 8.96� 10�7 Goethite 2.42
Fe2þ 4.70� 10�6 Magnetite 5.80
FeOHþ2 1.00� 10�8 Hematite 7.23
FeSO�4 1.55� 10�7 Maghemite �0.57
HSO�4 5.73� 10�7 Lepidocrocite 1.54
SO2�4 1.34� 10�4 H-Jarosite �4.37
2 Cl� 1.61� 10�4 Ferrihydrite 0.34
Fe ðOHÞþ2 1.93� 10�6 Goethite 3.04
Fe2þ 4.10� 10�5 Magnetite 8.57
FeOHþ2 1.11� 10�8 Hematite 8.48
FeSO�4 1.39� 10�6 Maghemite 0.68
HSO�4 3.03� 10�7 Lepidocrocite 2.16
SO2�4 1.39� 10�4 H-Jarosite �3.63
3 Cl� 1.68� 10�4 Ferrihydrite 1.27
Fe ðOHÞþ2 9.88� 10�6 Goethite 3.97
Fe2þ 7.09� 10�5 Magnetite 11.11
FeOHþ2 3.40� 10�8 Hematite 10.34
FeSO�4 2.46� 10�6 Maghemite 2.53
HSO�4 1.86� 10�7 Lepidocrocite 3.09
SO2�4 1.41� 10�4 H-Jarosite �1.70
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decreasing those of amorphous Al(OH)3, boehmite, and basaluminite.The solution from lysimeter 3 was undersaturated with respect tobasaluminite when DOC was considered.
CONCLUSIONS
The sampling and speciation methods used in this study alloweddetermination of the quantities of anionic and hydrophobic componentsof dissolved Fe, as well as those of reduced and total Fe in Aquodsolutions. The sum of anionic and hydrophobic Fe was interpreted to beorganically-complexed Fe, which was generally greater than 50% of totalFe in OA and Bhs horizon solutions. Surprisingly, organically-complexedFe was also a significant component of total Fe in the Bs horizon, andreduced Fe was often greater than 50% of total Fe in surface horizonsolutions. Iron reduction was clearly evident in the Bs horizon as the soiltemperature increased and the water level in the soil dropped. The soilsolids appear to be dominated by crystalline and amorphous forms of Fein the B horizons, with Fe solubility possibly controlled by ferrihydriteand maghemite. However, mineral equilibria calculations are performedwithout due regard to solution organic components, which may shiftequilibria in favor of Fe dissolution. This study provides evidence that amore detailed, field-based examination of the specific organic compoundsin soil solutions is worthwhile for improving Fe solubility calculationsin seasonally saturated soils.
ACKNOWLEDGMENTS
This is Scientific Contribution Number 2149 from the NewHampshire Agricultural Experiment Station.
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