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Sequential Selective Extraction Procedures for the Study of Heavy Metals in Soils, Sediments, and Waste Materials—a Critical Review

Amir Hass ab;Pinchas Fine c a USDA-ARS, Appalachian Farming Systems Research Center, Beaver, West Virginia, USA b Gus R. Douglass Land-Grant Institute, Agricultural and Environmental Research Station, West Virginia State University, Institute, West Virginia, USA c Ministry of Agriculture, Agricultural Research Organization, Volcani Center, Institute of Soil, Water and Environmental Sciences, Department of Soil Chemistry and Plant Nutrition, Bet-Dagan, Israel Online publication date: 27 April 2010 To cite this Article Hass, Amir andFine, Pinchas(2010) 'Sequential Selective Extraction Procedures for the Study of Heavy

Metals in Soils, Sediments, and Waste Materials—a Critical Review', Critical Reviews in Environmental Science and Technology, 40: 5, 365 — 399 To link to this Article: DOI: 10.1080/10643380802377992 URL: http://dx.doi.org/10.1080/10643380802377992

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Critical Reviews in Environmental Science and Technology, 40:365–399, 2010 Copyright © Taylor & Francis Group, LLC ISSN: 1064-3389 print / 1547-6537 online DOI: 10.1080/10643380802377992

Sequential Selective Extraction Procedures for the Study of Heavy Metals in Soils, Sediments, and Waste Materials—a Critical Review AMIR HASS1,2 and PINCHAS FINE3 Downloaded By: [US EPA Environmental Protection Agency] At: 20:16 14 May 2010

1

USDA-ARS, Appalachian Farming Systems Research Center, Beaver, West Virginia, USA Gus R. Douglass Land-Grant Institute, Agricultural and Environmental Research Station, West Virginia State University, Institute, West Virginia, USA 3 Ministry of Agriculture, Agricultural Research Organization, Volcani Center, Institute of Soil, Water and Environmental Sciences, Department of Soil Chemistry and Plant Nutrition, Bet-Dagan, Israel 2

The authors review selected protocols of sequential selective extraction procedure that are used to characterize the geochemical distribution of heavy metals in soils, wastes, and sediments. They discuss the development of earlier protocols, their modifications, and the extent to which a given protocol pertains to different conditions. Emphasis is given to the considerations that led to a choice of reagents for each step and to their order in the sequence. Published studies are used as case studies to critically evaluate the implied geochemical components of operationally defined extraction steps. Also assessed are possible effects of subsequent extraction steps and conditions on the selective dissolution of the solid components and their operational definitions. KEY WORDS: biosolids, iron oxide, metal partitioning, metal solubility, metal speciation

INTRODUCTION Selective sequential extraction procedures (SSE) are widely used in exploring the geochemistry of heavy metals in soils, sediments, and waste materials Address correspondence to Pinchas Fine, Ministry of Agriculture, Agricultural Research Organization, Volcani Center, Institute of Soil, Water and Environmental Sciences, Department of Soil Chemistry and Plant Nutrition, PO Box 6, Bet-Dagan 50250, Israel; E-mail: [email protected] 365

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(Beckett, 1989; Lake et al., 1984; Tessier & Campbell, 1991). The SSE method is based on the exposure of a sample to a sequence of wet chemistry assays. At each step of the sequence the sample is extracted in a solution that assumes to selectively remove or dissolve a specific form or solid component that the metal in question is associated with. The conditions in the extraction solution (e.g., presence of ligands or predetermined pH) are also set to minimize resorption and precipitation of the extracted entity. As the sequence advances, the reagents become more aggressive and their degree of specificity decreases. There are two main advantages in using SSE in evaluating metal status in soils and sediments. One is the ability to relate a specific extraction in the sequence to a specific chemical entity (e.g., carbonates, iron oxides) or form (e.g., soluble, exchangeable, sorbed) that the metal is associated with. The other is making inferences, based on such solid-phase speciation, on the behavior of the metal in the environment (e.g., mobility and bioavailability), or its response to changes in environmental conditions (e.g., solubility changes due to pH or Eh changes). Hence, the quality of an SSE protocol is in the ability of its reagents sequence to selectively and specifically extract, and to coherently define their respective intended phases. Even though the SSE concept for studying the behavior of heavy metals was developed in the 1970s (Gibbs, 1973; McLaren & Crawford, 1973a; Shuman, 1979; Stover et al., 1976; Tessier et al., 1979), a unified or agreed protocol has not yet been accepted. The number of steps and operationally defined phases in SSE schemes vary between three (Quevauviller et al., 1994; Silviera & Sommers, 1977) and nine (Krishnamurti & Naidu, 2000, 2002; Miller et al., 1986b) according to preferences and purpose. In general, the metal extraction steps are in the following order: Aqueous, and/or electrolyte solution to extract the soluble, and/or exchangeable metal species, respectively; a weak acid for metals bound to carbonates; a weak reducing agent to dissolve metals within Mn oxyhydroxides and amorphous iron oxyhydroxides; a weak oxidizing agent to characterize metals in the organic component; a strong reducing agent to characterize metals associated with well-crystallized iron oxyhydroxides, and a strong mineral acid(s) to dissolve the residue remaining after all of the preceding steps; mostly alumosilicates. Doubts over extraction specificity and selectivity promoted changes in both the choice of extractants and their order in the sequence. SSE protocols seldom achieve any of the following: (a) complete dissolution of a specific target component, (b) selective extraction or dissolution of the target component, and (c) full recovery of dissolved metals without their resorption or precipitation. It also led to rigorous criticism of SSE (e.g., Kheboian & Bauer, 1987; Nirel & Morel, 1990; Rendell et al., 1980). Several thorough reviews of SSE protocols have been published (e.g., Beckett, 1989; Gleyzes et al., 2002; Pickering, 1981; Sheppard et al., 1997; Shuman, 1991; Tessier & Campbell, 1991). Recently, Filgueiras et al. (2002)


Sequential Selective Extraction for the Study of Heavy Metals

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reviewed nearly 300 published studies on geochemical distribution of heavy metals in soils, sediments, and biosolids that used a wide selection of SSE protocols. They pointed out that the absence of a unified accepted protocol made it rather difficult to compare results from different studies. The present review encompasses most recent SSE derivatives as well as providing a systematic, chronological, and conceptual development of SSE protocols in the last four decades. It uniquely emphasizes and exemplifies the effect of choice of reagents, reaction conditions, and order in the sequence of an extraction step on the resulting interpretation of the operational solid-phase definitions. We examine some of the basic theoretical considerations that early SSE protocols were based on and highlight potential pitfalls that might occur if the underling principles and conditions are disobeyed or are applied to incompatible systems.

EARLY KEY STUDIES ON SSE Stover et al. (1976) developed an SSE protocol to characterize the distribution of metals in anaerobically stabilized sewage sludge. In this anaerobic system they assigned all acid (1 M HNO3 ) extractable metals to sulfides and considered as negligible the possible presence of oxyhydroxides (Table 1). Inasmuch as Stover et al. (1976) considered not including a discrete extraction step for oxides might be justified for an anaerobic system, Schalscha et al. (1980) extended the use of this protocol to characterize metals in soil irrigated with sewage wastewater. Later, Emmerich et al. (1982) and Sposito et al. (1982; followed by Alva et al., 2000; Chang et al., 1984; Sims & Kline, 1991; Tsadilas et al., 1995) used the same protocol as the basis for their examination of the distribution of metals in soils with applied sewage sludge. In these studies, the sulfide component was considered part of the residual component, and the adsorbed component was extracted with distilled water (instead of NaF; Sposito et al., 1982; Table 1). Because the adsorbed metal content was relatively low, this component was often combined with the exchangeable component (Alva et al., 2000; Sims & Kline, 1991; Sposito et al., 1982). Applying an SSE scheme that was designed for a strictly anaerobic medium to an aerobic soil system is problematic for at least two reasons: for (a) not selectively extracting the free oxide component, and (b) incorrectly identifying oxyhydroxides-associated metals as sulfides. Thus, oxide-bound metals interpreted as sulfide-bound metals might be regarded as a potential source under oxidizing conditions rather than as a stable entity and a sink. The previously mentioned studies, which used Stover et al.’s (1976) extraction step for sulfides (1 M HNO3 ; Table 1) are puzzling because the reagent was shown to merely partially dissolve metal sulfides (Rudd et al., 1988) even in anaerobic sludge samples (Hullebusch et al., 2005). Rudd

368

Soluble

McLaren and Crawford, 1973 Gupta and H2 Oa Chen, 1975 Stover et al., 1976 Tessier et al., 1979 Shuman, 1979 Sposito et al., 1982 Salomons and Forstner, 1984 Shuman, 1985 Miller et al., H2 Oa 1986b McGrath and Cegarra, 1992 Ma and H2 Oa Uren, 1995

Gibbs, 1973

Source

NH4 OAc. EDTA, pH 8.3b

CaCl2

HOAc., pH 2.5d Na2 H2 EDTA

NaOAc, pH 5b

EDTA

EDTA pH 6.5 NaOAc. pH 5b

HOAc.c

Carbonate

H2 O2 , NaOAc., pH 4.74e pH 4.74d

NaOH

NaOCl, pH 8.5 K4 P2 O7 f

Mg(NO3 )2 , pH 7 Ca(NO3 )2 b Pb(NO3 )2 c

H2 O2 , pH 2e

NH4 OAc, pH 7a

NaOH

KNO3 H2 O

H2 O2 , pH 2d H2 O2

MgCl2 , pH 7a MgCl2 pH 7

Na4 P2 O7

KF, pH 6.5

KNO3

NaOCl pH 8.5c K 4 P2 O7

Organic matter

H2 O2 e

HOAc.

Adsorb

NH4 OAc.b

MgCl2 a pH 7 CaCl2

Exchangeable

NH4 Ox. in the NH4 Ox Ascorbic dark, pH 3 acid, pH 3 NH4 Ox. in the NH4 Ox.Under g dark, pH 3 UV, pH 3h

NH4 Ox. HOx., pH 3d

NH4 Ox. in the dark, pH 3

HA, HOAc., pH2c

HA, HOAcf

NH4 Ox. UV, pH 3.25

Quinon, NH4 Ox., pH 3.25f NH4 OAc., pH 7c

HA, pH 2e

HA, pH 2

HA, pH 2c

b

Crystalline iron oxides

Citrate NaDithionite

Amorphous iron oxides

HAd, pH 2

Mn oxides

Operationally defined component

1M HNO3

Sulfides

H2 SO4 ·HClO4 · NO3 —sumg

HCl·HNO3

HCl·HNO3 ·HFi

HCl·HNO3 ·HF

HNO3 f

4M HNO3

HCl·HNO3 ·HF

HCl·HF·HClOe

HCl·HF·HclO4 g

LiBO2 (1000◦ C)· HNO3 d HF

Residual

TABLE 1. Selective sequential extraction—selected protocols (extraction solutions, sequences, and operationally defined components)


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NaOAc. EDTA, pH 4.65b

NaOAc., pH 4.74e

EDTA pH 6.5

NaOAc., pH 5b EDTA

HOAc., pH 2.6c

Na4 P2 O7 e H2 O2 , pH 2e Na4 P2 O7

NaOAc., pH 5b

H2 O2 , pH 2d

NH4 F pH 6.5 Na4 P2 O7 H2 O2 , pH 2c NH4 OAc. H2 O2 , pH EDTA, pH 4.74f 8.3c HOAc., H2 O2 , pH pH 5b 2d § Na4 P2 O7 NaOAc., NaOCl, pH 5 pH 8.5 Na4 P2 O7 d

Pb(NO3 )2 NH4 OAc., pH 6b

HOAc. Ca(NO3 )2 c

Na4 P2 O7 e

HA, pH 2b

NH4 Ox in the dark, pH 3f HA HOAc. pH 2d

NH4 Ox. Ascorbic acid, pH 3.1g

NH4 Ox. in the NH4 Ox. HOx dark, pH 3f under UV, pH 3g HA, HOAc., pH 2e

Quinon, NH4 Ox., pH 3.25g NH4 OAc., pH 7d HA, pH 2c NH4 Ox. in the NH4 Ox. Ascorbic dark, pH acid, pH 3f e 3.0 HA, pH 2 NH4 Ox. in the 6M HCl dark, pH 3.0 NH4 Ox., pH 3.25c

HA, pH 2c

HA, HOAc., pH 2c HA, pH 2d

HA pH 2d

6M HNO3

7M HNO3 – sume

HNO3 ·HCl

HF·HClO4 g

3M HNO3 HF·HNO3 , HOAc.d H2 SO4 ·HClO4 · NO3 – sumh

18M HNO3

4M HNO3 f

H2 SO4 ·HClO4 · HNO3 h

4M HNO3 f

HNO3 ·HCl·HFh

Note. Alphabetic superscripts (a, b, c, d, e, f, g, h) indicate the sequence order when differing from the order of the table columns. Na4 P2 O7 was used as a seperate step, following the specifically adsorb, to characterize metal-organic complex bond.HA = hydroxylamine hydrochloride; OAc. = acetate; OX = oxalate; ETDA = ethylendiamine tetra acetic acid.

Krishnamurti NH4 NO3 , pH 7a and Naidu, 2000 Silveira et CaCl2 al., 2006 Fitamo et al., NH4 NO3 , pH 7a 2007

Sloan et al., Ca(NO3 )2 a 1997 Smith et al., CaCl2 1997 Li et al., 1997 H2 O KNO3 Nyamangara, HOAc.a 1998 Ma and H2 Oa MgCl2 , Uren, 1998 pH 7b

Berti and H2 Oa Ca(NO3 )2 , Jacobs, pH 7b 1996 Han and NH4 NO3 , Banin, pH 7a 1996 Iwasaki NH4 OAc., et al., 1997 pH 7a



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et al. (1988) demonstrated incomplete extraction of metal sulfides (especially those of Zn and Cu) in 6 M HNO3 extraction, either in a single step or in a sequential extraction. Shuman (1979), and McLaren and Crawford (1973a), working on acid and neutral soils (4.5 < pH < 7 and 3.2 < pH < 7.3, respectively), did not extract the carbonate component separately, and extracted all oxyhydroxides in one step (Table 1). Shuman (1979) also determined metals in the residual component of the sand, silt, and clay size fractions separately. Later, Shuman (1985) revised his former protocol, introducing extraction steps for manganese oxyhydroxides and amorphous iron oxyhydroxides, and for well-crystallized iron oxyhydroxides (Table 1). To suppress manganese oxyhydroxides dissolution in the preceding organic matter extraction step, Shuman (1985) extracted the organic matter in an alkaline medium. Li and Shuman (1996) and Chowdhury et al. (1997) used the later protocol in treating acid soils (pH 4.4–5.6 and < 6.4, respectively). Neilsen et al. (1986) used the Shuman (1979) protocol to study soils encompassing a wide range of pH values (4.05–8.25). Applying this protocol to neutral and alkaline soils without addressing metals possibly associated with carbonates and other acid-sparingly dissolved minerals (e.g., calcium phosphates) is likely to result in overestimating the metals associated with the organic matter (if following Shuman, 1979) or with manganese oxyhydroxides (if following Shuman, 1985). Tessier et al. (1979) developed an SSE scheme designed to characterize metals by their pollution potential (Table 1). This protocol was widely used as is (e.g., Chlopecka, 1996, Chlopecka & Adriano, 1996; Dollar et al., 2001; McLaren & Clucas, 2001; Wang et al., 1997) or with adaptations (e.g., Banin et al., 1990, Brennan, 1991; Kryc et al., 2003; Ma & Uren, 1995, 1997; Salomons & Frostner, 1984). Salomons and Forstner (1984) split Tessier et al.’s (1979) oxyhydroxides extraction into two steps: a milder reduction for manganese oxyhydroxides and amorphous iron oxyhydroxides, followed by a strong reduction of the well-crystallized iron oxyhydroxides (Table 1). Hullebusch et al. (2005), largely adopted the Tessier et al. (1979) scheme in their study of metals in anaerobically digested sludges, but skipped the oxides step, following Stover et al. (1976) in this respect. Peltier et al. (2005) chose to apply the Tessier et al. (1979) protocol in full to study anoxic wetland sediments. They did point out that metal sulfides could be mistakenly identified as metal oxides. Tessier et al. (1979) noted that specifically adsorbed metals were included in the carbonate step. This was treated by Ma and Uren (1995) by adding a chelating agent (ethylendiamine tetra acetic acid [EDTA]) in the exchangeable component extraction step to also remove specifically adsorbed metals prior to carbonates dissolution (Table 1). In a later protocol, Ma and Uren (1998; Table 1) extracted the specifically adsorbed component in a separate step. Rudd et al. (1988) adapted the exchangeable (1 M



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MgCl2 ) step from Tessier et al.’s (1979) protocol in their concise SSE protocol. They adjusted the pH of the exchangeable step to 5 (rather then to 7) to avoid the concern of removing specifically adsorb metal in the proceeding acidic NaOAc step (Rudd et al., 1988). The inclusion of sorbed metals in the carbonate step of Tessier et al.’s (1979) protocol was circumvented by simply changing the operational definition from the carbonates component to specifically sorbed–carbonate bound (Ahnstrom & Parker, 1999), specifically sorbed (McLaren & Clucas, 2001) or acid extractable (Wong & Selvam, 2006). SSE protocols clearly have many potential pitfalls, and the abundance of different protocols make it difficult to make comparisons between studies. In an attempt to provide an internationally accepted SSE protocol, the European Commission’s Bureau Community of Reference (BCR) developed a simplified three-step SSE protocol (Quevauviller, 1998a; Quevauviller et al., 1994). The BCR also established standard reference materials for SSE quality control (Quevauviller, 1998b). The first step in the protocol is extraction with 0.11 M HOAc at 20◦ C overnight, the second step treats the residue with 0.1 M hydroxylamine hydrochloride (NH2 OH-HCl) at pH 2 and 20◦ C overnight, and the third step is treatment with 30% H2 O2 at pH 2–3 and 85◦ C for 1 hr, followed by 1 M NH4 OAc at pH 2 and 20◦ C. The second step was later amended by raising the concentration of NH2 OH-HCl to 0.5 M, and by adding more acid (Rauret et al., 1999). In addition, a fourth step was added to satisfy the International Standards Organization (ISO) requirement for an internal check, by digesting the last residue in aqua regia. The sum of metal extracted in the four extraction steps was thus compared with metal content obtained from digestion of an intact sample in aqua regia (Rauret et al., 1999). The BCR protocol steps were merely numbered with no attempt to imply specific chemical entities (Rauret, 1998; Rauret et al., 1999). Nyamangara (1998; as well as Basta et al., 2005; Kartal et al., 2006) defined the four steps of the modified BCR protocol sequence as exchangeable, Fe-Mn oxides, organically bound, and residual. Hullebusch et al. (2005) referred to the first step as exchangeable + water, and acid soluble. Filgueiras et al. (2002) categorized the second step of the BCR SSE as Mn-oxides and the third step as oxidizable oxides and sulfides. By limiting the number of extraction steps, metal extraction in each step was increased. Inasmuch as this promoted higher degree of repeatability and reproducibility, it did not make the protocol more specific and selective with respect to intended phases of interest. Hence, the ambiguity about the operational definition of the components, which triggered the EC-BCR pursuit of a standardized SSE protocol, was not reduced. It should be noted that the residual was added to the BCR protocol merely to satisfy ISO requirements for an internal check (Rauret, 1998; Rauret et al., 1999). It was neither added as a final SSE step nor does it correspond to the residual component of other SSE protocols. Obviously, the residual component reflects the thoroughness of the preceding steps. In other SSE protocols, the residual component is

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largely referred to as metals residing within the crystal lattice of refractory minerals. However, in the BCR protocol, incomplete dissolution of crystalline iron oxides under the conditions of the second step is expected to contribute to the fourth step as well as to the third step (as also noted by Rauret et al., 1999). Other shortcomings of the BCR protocol, including the nonselective nature of the first step (e.g., acetate and pH effect on Mn oxide dissolution), as well as the nonspecific nature of the second step (e.g., incomplete dissolution of iron oxides by hydroxylamine), and their effect on extractable metals in subsequent steps, are discussed subsequently.


STRUCTURE OF COMMON SSE PROTOCOLS The Soluble Component Extraction with distilled water by shaking for 16 hr is used to characterize soluble metals in soils. This is the first step in the sequence. Because of the relative low content of soluble metals, sometime below the detection limit, this component is often omitted, and the soluble form is included in the next step (Table 1). Water extraction may have side effects, including solubilizing organic matter, the effect of which is discussed in the exchangeable component step.

The Exchangeable Component This fraction is attributed to metals that are associated with solids by electrostatic exchange mechanisms. The extraction is done by shaking the sample in the presence of a concentrated solution of an inert electrolyte (Tessier & Campbell, 1991). The electrolytes commonly used are nitrate or chloride salts of ammonium, calcium, potassium, or magnesium (Table 1). Because chloride ions tend to form soluble complexes with some metals (e.g., cadmium; Bingham et al., 1984; Doner, 1978; Lumsdon et al., 1995), the use of nitrate salts in this step is preferred (Beckett, 1989). Extraction using NH4 OAc at pH 7, which is a standard method for the characterization of exchangeable cations in soils (Thomas, 1982), has also been used in SSE protocols (Brennan, 1991; Brown et al., 1984; Iwasaki et al., 1997; Salomons & Forstner, 1984). Nyamangara (1998) who used the modified BCR SSE define this component (i.e., exchangeable) by extraction with acetic acid, the first step in the BCR SSE (0.11 M HOAc; Table 1). Selectivity was severely affected when acetic acid or even acetate salts were used to extract this component. Tessier et al. (1979) found that 1 M NaOAc dissolved carbonates even at pH 8.2, despite the fact that this pH is employed to prevent carbonate dissolution (and their contribution to the exchangeable cations; Bower et al., 1952). Gibbs (1973) found that NH4 OAc dissolved oxyhydroxide coatings of riverbed sediments. Stover et al. (1976),



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Miller et al. (1986a) and Berti and Jacobs (1996) reported lack of selectivity of HOAc because of its ability to dissolve a wide range of soil minerals. This is further discussed subsequently with respect to the specifically adsorbed component. This extraction step is performed at a fixed pH (usually 7) or without pH adjustment (Table 1). Any adjustment of soil pH can considerably influence the amount of metals extracted. This is due to pH effects on mineral solubility, mineral surface charge (especially in variable charged soils), metal speciation in solution (metal hydroxylation), and organic matter solubility and configuration. Neilsen et al. (1986) compared extraction of exchangeable Zn from a range of soils using either 1 M MgCl2 solution without pH adjustment, or 1 M NH4 OAc solution at pH 7. In the acidic soils tested, less Zn was extracted with NH4 OAc than with MgCl2 , and the difference increased with an increase in soil acidity. Neilsen et al. (1986) suggested that more specific adsorption or precipitation of Zn had occurred during the extraction at a pH higher than the native soil pH. They concluded that characterization of the exchangeable component should be done at the indigenous soil pH. Naidu and Harter (1998) examined the ability of different ligands (nitrate, acetate, oxalate, and citrate) to recover cadmium salts applied to different soils. They found that the recovery decreased as the pH of the extraction solution increased. Dyer et al. (2004) in a Zn-ferrihydrite system showed that the apparent log K for Zn adsorption by ferrihydrite increases linearly with an increase in pH (from 2.18 at pH 4 to 10.32 at pH 8). Cottenie and Kiekens (1972) illustrated the effect of the pH of the 1 M NH4 OAc extractant on the recovery of previously sorbed metals. In a batch experiment, they equilibrated a soil (pHw 6.3, CEC 11.5 cmolc kg−1, 22.3 g C kg−1) with solutions of increasing concentrations (0–50 mg kg−1) of Fe, Mn, Zn, or Cu, and at pH values that ranged from the native soil pH down to pH 0.5. After a 30-min sorption period the excess metal was removed by washing with ethanol. Metal-loaded soil samples were then subject to extraction with 1 M NH4 OAc at either pH 7 or 3. They recorded the added metal that remained in solution (Curve 1 in Figures 1 and 2), was retained by the soil (Curve 2 in Figures 1 and 2), and was recovered by 1 M NH4 OAc extraction at pH 7 and at pH 3 (Curves 3 and 4, respectively, in Figures 1 and 2). They found that the pH of the desorbing solution strongly affected the recoveries of Zn and Cu but not those of Mn and Fe. For example, at a 20 mg kg−1 loading rate, all the Cu and Zn that were retained by the soil at an equilibrating pH > 4 (Curve 2 in Figures 1 and 2) were fully recovered by 1 M NH4 OAc at pH 3 (Curve 4 in Figures 1 and 2), but not by 1 M NH4 OAc at pH 7 (Curve 3 in Figures 1 and 2). They attributed the unrecovered metals to fixed metals. It is noteworthy that the effect persisted at an adsorption equilibrium pH in which metal precipitation as oxides is insignificant (at pH < 5.5). It is possible that changes in metal–surface interactions, from outer- (exchangeable) into inner-sphere complexes (specifically adsorb), or


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FIGURE 1 Displacement of adsorbed Zn by NH4 OAc. Soil was equilibrated with 20 mg kg−1 Zn at different pH levels for 30 min (equilibrium pH is indicated as pH of filtrate). Equilibrium solution was then displaced and the soil was washed with ethanol prior to NH4 OAc extraction. Curve 1 = added Zn remained in equilibrium solution; Curve 2 = added Zn retained by the soil after equilibration and prior to extraction; Curve 3 = NH4 OAc pH 7 extractable Zn; Curve 4 = NH4 OAc pH 3 extractable Zn (Cottenie & Kiekens, 1972; with permission of the International Potash Institute [IPI], Horgen, Switzerland).

precipitation as amorphous oxides during the extraction at pH 7 could have occurred. The effects of pH on soil organic matter solubility was used by Impellitteri et al. (2002) to explain the pH dependency of Cu, Ni, Cd, and Pb extractability in deionized water from 18 Dutch soils. The metal–pH relationships were rather parabolic, with a minima near pH 5. The magnitude of the metal solubility increase at pH > 5 was metal-specific: it was highest for Cu (especially at pH > 7, which also exceeded the Cu solubility at the low pH range) and lowest for Zn, which did not increase at pH > 5. They also noticed that the solubility of soil organic matter, and that of the humic



375

FIGURE 2 Displacement of adsorbed Cu by NH4 OAc. Soil was equilibrated with 20 mg kg−1 Cu at different pH levels for 30 min (equilibrium pH is indicated as pH of filtrate). Equilibrium solution was then displaced and the remaining soil was washed with ethanol prior to the extraction. Curve 1 = added Cu remained in equilibrium solution; Curve 2 = added Cu retained by the soil after equilibration and prior to extraction; Curve 3 = NH4 OAc pH 7 extractable Cu; Curve 4 = NH4 OAc pH 3 extractable Cu (Cottenie & Kiekens, 1972; with permission of the International Potash Institute [IPI], Horgen, Switzerland).

acid component of the dissolved organic carbon (DOC), both increased with increasing extraction solution pH. Hence, they concluded that the increase in metal solubilization at pH levels greater than 5 resulted from an increase in DOC and humic acid solubility, and complexation of the metals with the dissolve organic matter. Similar parabolic behavior of Cu solubility (with minima at pH 5–7) and an increase in soluble Zn with a decrease in pH was noted by He et al. (2006) when leaching cultivated sandy soil (pH 7.0) with deionized water at different pH levels. This mode of behavior was less pronounced in an adjacent acid forest soil (pH 4.2) in which the extraction solution pH effect (pH 3–9) on Cu solubility was minimal. You et al., (1999), working on 15 New Jersey soils (mostly Ultisols and Alfisols, pH 4.2–6.4), also found an increase in soluble organic matter and the proportion of


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FIGURE 3 Dissolution of soil organic matter by 0.01M NaNO3 soil extract at different pH. HA = humic acid; FA = fulvic acid (reprinted from Sci. Total Environ., 227, 155–160, You et al., Partitioning of organic matter in soils, 1999, with permission from Elsevier).

humic acid in the dissolved organic matter (DOM), with an increase in the pH of the 0.01 M NaNO3 extraction solution (from 3.9 to 7.5; Figure 3). It should be noted that an increase in the proportion of soluble humic acid with an increase in pH results in a more stable DOM–metal complex, as the humic acid metal complex is more stable than the fulvic acid metal complex (Stevenson, 1994). Ionic strength and composition may also have a substantial effect on metal solubility by influencing the dispersion and flocculation of soil particles. Miller et al. (1986a) found that deionized water extracted more Cu than 0.5 M Ca(NO3 )2 or 0.05 M Pb(NO3 )2 solutions. Moreover, a 0.025 M Ca(NO3 )2 solution extracted more Cu than did a 0.5 M solution of the same salt. The authors suggested that at the higher electrolyte concentration the DOC flocculated and precipitated together with metals that were associated with it, resulting in an apparent decrease in metal solubility. Similar effects on organic matter solubility were reported by Reemtsma et al. (1999). They extracted twice as much DOC from a loamy sand surface soil of a wastewater infiltration site using deionized water than when using 0.05 M CaCl2 (86 and 44 mg DOC kg−1, respectively). Similarly, Oste et al. (2002) demonstrated the effect of pH and Ca concentration (added as Ca[NO3 ]2 from 0.00 to 0.025 mol kg−1 dry soil) on DOC solubility in nine agricultural soils (pH 4.2–6.9). DOC solubility increased with increasing pH (range = 6.0–8.7) and decreased with an increase of Ca in solution. Ghosh and Schnitzer (1980) showed that samples of humic and fulvic substances extracted from soils were flexible linear colloids at low ionic strength (and also at low sample concentration and neutral or high pH). The extended configuration was due to charge separation and intramolecular repulsion. As the ionic strength increased (≥0.05 M NaCl), these



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macromolecules spontaneously assumed coiled configurations similar to those of rigid uncharged spherocolloids. Similarly, Tsutsuki and Kuwatsuka (1984) demonstrated a reduction in the effective mean size distribution of humic materials that were extracted from a soil. The effective size decreased from 210 and 310 kDalton to 2.3 and 4.5 kDalton, respectively, following an increase of the ionic strength from 0.03 to 0.5 M NaCl. The quantity of metals that can be extracted is influenced by the following: the tendency of a given metal to interact with soil DOC, the properties of the DOC, the composition of the mineral phase and its interaction with the DOC, and the tendency of the dissolved complexes to flocculate and disperse in response to the extraction conditions (e.g., solid: solution ratio, pH, ionic strength, ionic composition). For example, Romkens and Dolfing (1998) studied the effect of Ca addition on the solubility of previously extracted DOC (using 0.5 N NaNO3 , pH 7.5) from manure amended, and from inorganic fertilized agricultural sandy soils (pHKCl 4.7–5.5). They found that the addition of 10 mM CaCl2 to the extracts reduced DOC and soluble Cu concentrations by nearly 50%. Most of the flocculated DOC was of the high molecular weight fraction (Romkens & Dolfing, 1998). He et al. (2006) compared amounts of extractable Cu and Zn in 13 cultivated and forest soils in Florida (pH 4.12–7.6). Extraction was done with deionized water or with different nitrate salts (i.e., NH4 NO3 , KNO3 , NaNO3 , and Ca[NO3 ]2 ), all at 0.02 N. Sodium nitrate and deionized water extracted the most Cu, whereas Ca(NO3 )2 extracted the least. In the case of Zn, deionized water and Ca(NO3 )2 extracted more Zn than did the other extractants. They suggested that ionic strength and salt composition affected DOC solubility and thus Cu extractability, whereas Zn solubility depended more on exchange reactions. Assuming that no specifically adsorbed ions, ligands, or potential determining ions (that might induce surface charge and potential changes) are introduced at this step, and taking into account colloid stability considerations (i.e., the Schulze-Hardy rule and the DLVO theory; Buffle et al., 1998; Sposito 1984), coagulation effects of the extraction solution are expected to increase with an increase in ionic strength, ion charge, and charge density, and with an increase in solid–solution ratio. Hence, polyvalent ions are more effective coagulants than monovalent ones, and Ca, for example, is a more efficient than Mg. Given the previous explanation, the use of Ca salts to evaluate contaminants extractability, such as using 0.5M Ca(NO3 )2 in the Potentially Bioavailable Assessment Sequential Extraction procedure (PBASE; Basta et al., 2001) or in SSE to characterize soluble and exchangeable form (e.g., 0.5M Ca[NO3 ]2 by Sloan et al., 1997; 0.1 M CaCl2 by Silveira et al., 2006; Table 1), and especially in organic-waste amended soils, could result in coagulation and precipitation of soil colloids and dissolved organic-metal complexes.


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FIGURE 4 Adsorption envelope of cadmium on goethite at different ionic strength (reprinted from J. Colloid Inferf. Sci., 115, 564–572, Hayes and Leckie, Modeling ionic strength effects on cation adsorption at hydrous oxide/solution interfaces, 1987, with permission from Elsevier).

The Specifically Adsorbed Component This component, attributed to metals in direct chemical association with surfaces (i.e., inner-sphere complex; Sposito, 1984), does not always receive a separate extraction step in SSE protocols (Table 1). An important macroscopic characteristic of specific adsorption of metals is its independence on solution ionic strength (Figure 4; Hayes & Leckie, 1987). Swallow et al. (1980) reported that Cu and Pb sorption onto hydrous ferric oxide was unaffected by the ionic strength of the background solution (0.005–0.5 M) or by its composition (NaClO4 or complex artificial seawater mix). The only exception was Pb at high Cl concentrations, attributed to Pb-Cl complexation. Chibowski and Janusz (2002) showed no effect of an increase in NaClO4 background solution (10−3–10−1 M) on Zn adsorption on hematite. Dyer et al. (2004) showed similar behavior for Zn adsorption on ferrihydrite in NaNO3 background solutions of the same concentrations. Kinniburgh et al. (1976) reported specific adsorption of heavy metals by iron and aluminum oxyhydroxides in background solutions of 1 M NaNO3 . Millward and Moore (1982) tested Cu, Mn, and Zn adsorption to iron oxyhydroxides from model seawater at concentrations up to 0.42 M Cl. Increasing salt concentrations somewhat decreased Cu adsorption but not that of the other metals. Extraction of this component is based on competitive adsorption using ligands that compete with either the metals for adsorption sites on mineral surfaces or the surface to form soluble complexes with the metals, or a combination of the two. Hence, Stover et al. (1976) used KF to characterize



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the adsorbed fraction (of Cd, Cu, Ni, Pb, and Zn); Miller et al. (1986a) used lead as a displacing ion to characterize specifically adsorbed copper; and Iwasaki et al. (1997) used lead (Pb[NO3 ]2 ) in ammonium acetate solution to characterize adsorbed copper and manganese. Ma and Uren (1998) used a chelating agent (EDTA) in NH4 OAc solution for this purpose. In a number of studies (McLaren & Crawford, 1973a; Nyamangara, 1998; Payne et al., 1988; Silveira et al., 2006) acetic acid (at 0.11–0.44 M HOAc or as 1 M Na or NH4 OAc; Table 1) was used to extract specifically, as well as nonspecifically, adsorbed metals (NH4, Na, H+ serve as a displacing ion and acetate as a complexing agent). However, other studies showed that HOAc was not selective enough. Stover et al. (1976) reported the dissolution of carbonates and sulfides in 2.5% HOAc. Berti and Jacobs (1996) referred to the metals they extracted in 0.44 M HOAc (following extraction of the soluble and exchangeable components) as the acid soluble constituent, including carbonates and phosphates (Table 1). Miller et al. (1986a) reported the dissolution of copper hydroxide, copper hydroxycarbonate, and copper sulfide with 0.44 M HOAc (at pH 2.5). Miller et al. (1986b) developed a SSE for Cu in acid soils (5.3 < pH < 5.6). They concluded that 0.05 M Pb(NO3 )2 + 0.1 M Ca(NO3 )2 was more specific than 0.44 M HOAc (at pH 2.5) and thus selected it to extract specifically adsorb Cu (Miller et al., 1986b). Acetic acid was the next step in the sequence, and was used to extract the acid soluble Cu component (Table 1). Similarly, based on previous reports on the possible dissolution of carbonates in 1 M NH4 OAc at pH 7, but not in 1 M NaOAc at pH 8.2, Tessier et al. (1979) evaluated the latter in their search for an appropriate extractant for exchangeable metals. Nevertheless, the acetate dissolved carbonates, even in the sodium salt state and at the high pH (8.2). This led Tessier et al. (1979) to prefer the use of 1 M MgCl2 at pH 7 for this step in their protocol (Table 1). Considering extraction of a heterogeneous sample, the redistribution of metals from exchange sites to specific sorption sites is likely to occur under the conditions of the previous step (i.e., exchangeable). Furthermore, under exchangeable extraction conditions, soil homogenization (by grinding and sieving), and extraction procedure harmonization (i.e., shaking at high solution–soil ratio) may promote specific adsorption of electrostatically adsorbed metals. Homogenization and harmonization procedures as well as the extraction conditions and solution composition contradict the basic purpose of the SSE protocol, which is to provide specific and selective extraction conditions at each step. If differentiation between exchangeable and specifically adsorbed metals is a major objective, then the resorption phenomena, which is emphasized in subsequent steps needs to receive some attention here and in the preceding step in order to prevent resorption of released metal from exchangeable sites onto specifically adsorbed ones. This may also have some relevance in the initial step, when attempting to characterize (specifically and selectively) water-soluble metals.

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The Carbonate Component Carbonates dissolution is usually the first acid-extraction step in SSE sequences (Table 1). In studies in which extraction of the organic component is performed in an acid medium, the carbonate-extraction step precedes it (Han & Banin, 1997; Tessier et al., 1979). When the extraction of the organic component is performed in an alkaline medium the carbonate extraction step succeeds it (Stover et al., 1976; Sposito et al., 1982; Shuman, 1985). Ma and Uren (1995, 1998) extracted the carbonates after extracting manganese oxyhydroxides in a neutral medium (Table 1). Acid dissolution in 1 M Ammonium Acetate (NaOAc) at pH 5 or by ligand exchange in 0.1 M EDTA at pH 6.5 is often used to characterize metal carbonates (Table 1). The acetate and the EDTA are used also to stabilize the dissolved metals as soluble complexes, which inhibit possible resorption or precipitation. Using EDTA to dissolve carbonates has two main inherent difficulties. The first is that it may dissolve organic matter. Stover et al. (1976) overcame this problem by extracting the organic matter before extracting the carbonates (Table 1). The other difficulty is that EDTA dissolves iron oxyhydroxides (Borggard, 1988; Loveland, 1988). Hence, EDTA is seldom used (e.g., Fitamo, et al., 2007) when iron oxides are extracted in a subsequent discrete step (Table 1). Han and Banin (1995) studied the dissolution of soil carbonates in 1 M NaOAc at 1:25 solid–solution ratio. They found that at soil carbonate levels ≤ 300 g kg−1, buffering the extractant at pH 5.5 allowed complete carbonate dissolution. At 300–500 g carbonates kg−1, pH 5 was needed for complete dissolution, and at ≥ 500 g carbonates kg−1, the latter extraction had to be repeated twice for complete carbonate dissolution.

SELECTIVITY

PROBLEMS

Hanahan (2004) recently drew attention to selectivity problems with 1 M NaOAc (pH 5), demonstrating its ability to dissolve portlandite (Ca[OH]2 ) and brucite (Mg3 [OH]6 ). Earlier, Tessier et al. (1979) pointed out that the marked pH change in the carbonate dissolution step (i.e., from 7 in the preceding step to 5; Table 1) could result in metal extraction from other components too. This 2 pH unit reduction is at the critical pH range that covers the pH adsorption edge of most heavy metals of interest and of most relevant soil components (e.g., Zn on Mn oxide, Loganathan et al., 1977; Cd on soil, Elliott & Denneny, 1982; Cd, Cu, and Zn on acid soil, Spathariotis & Kallianou, 2001; Zn and Cd on Hematite, Chibowski & Janusz, 2002; Zn on ferrihydrite, Dyer et al., 2004; Figure 4). The proportion of metal that is specifically adsorbed, or in a broader term, that is retained in a nonexchangable form (i.e., operationally define as not extractable with 1 M NH4 OAc pH 7) increases



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sharply with an increase in pH in this range. Unless previously removed, the carbonate-bound metals step also represents specifically sorbed metals, contributed from all of the sample solid components. Furthermore, dissolution of Mn oxides in the 1 M NaOAc, pH 5 buffer can also occur. Han and Banin (1995) report a marked increase in V, Co, Ni, Cu, and Cr extracted at pH 5 compared to extraction at pH 5.5, in soils with carbonate content of 129–137 g kg−1. Although this increase in metal concentrations at the lower pH was accompanied by a marked increase in Mn, extractable Ca was the same at both pHs, and extraction at pH 5.5 was sufficient to assure complete carbonate dissolution. Tessier et al. (1979) reported high levels of extractable Fe and Mn in the NaOAc pH 5 buffer, which they attributed to dissolution of manganous and ferrous carbonates. The exceedingly high levels of Cu, Ni, Pb, and Zn found by Tessier et al. (1979) in this step, when comparing sedimentary rock-, deep-sea-sediment-, or biogenic-calcium carbonate, were referred to as specifically adsorbed metals that had been released by the pH change of this step. These metals exhibit high affinity towards Mn and iron oxides (e.g., Brummer et al., 1983; McBride, 1989; McKenzie, 1980; McLaren & Crawford, 1973b; McLaren et al., 1986). The pH range of 5–6 in general, and of pH 5.5 in particular, seem critical for Mn oxyhydroxide dissolution as well as Mn phytoavailability and toxicity (Bromfield, 1978; Bromfield & David, 1978; LaZerte & Burling, 1990; Sarkar & Jones, 1982; Sims, 1986; Sims & Kline,1991). Although Han and Banin (1995) suggested the use of pH 5 repeatedly, in carbonate rich soils, they reported that a pH lower than 5.8–5.9 in the supernatant at the end of the extraction was sufficient to assure complete carbonate dissolution. Owing to the susceptibility of Mn compounds to dissolution at lower pH levels and considering the results of Han and Banin’s (1995) study, it seems appropriate to set the pH of this step at 5.5 while checking that the final solution pH, by the end of the reaction, does not exceed 5.9.

OPERATIONAL DEFINITION Whether the extractable metal at this step results in overestimation depends also on the operational definition given to this step. Hence, this component was termed carbonate-bound (Han & Banin, 1997), acid extractable (Wong & Selvam, 2006), specifically sorbed (McLaren & Clucas, 2001), or specifically sorbed–carbonate bound (Ahnstrom & Parker, 1999). Hall et al. (1996) used NaOAc at pH 5 as the first extraction step with no previous water or salt extraction and termed it exchangeable, adsorb, and carbonates-bound metals. Such broadening of definitions obscures the chemical nature of the extracted metals. This may affect interpretations regarding processes that can take place in the medium studied. For example, decomposition of organic matter (especially in biowastes amended mineral soils) reduces the specific adsorbtion capacity of the soil, which may result in release of metals to the soil solution (McBride, 1995). However, this can also result in coprecipitation


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of the released metals as carbonates or phosphates solid solutions. Metals in a solid solution display very low solubility, depending on their activity in the solid (Papadopoulos & Rowell, 1988, 1989). Thus, metals can change over time from more labile, specific adsorbed sites on organic matter into insoluble inorganic forms (Basta et al., 2005; Logan & Chaney, 1983). The unified extraction step (i.e., the extraction of both specifically surface adsorbed components and the carbonates) would not enable researchers to distinguish between the two components and, hence, between the previously mentioned processes and redistributions. Hanahan (2004) demonstrated possible contributions of chemical entities other than carbonates to Tessier et al.’s (1979) carbonate component. He employed a modified Tessier SSE protocol on minerals (hematite, gibbsite, portlandite, calcite, and brucite) that were previously spiked with Cu or Pb by equilibrating in metal solutions for 24 hr. Metals recovered in the first step (MgCl2 , pH 7) accounted for < 1% of amounts added. Copper recovery in the carbonate step (Step 2) exceeded that in all the other steps combined, regardless of the host mineral. Lead recoveries in the carbonate step were lower only for hematite and brucite. Hanahan (2004) suggested that in some cases the recovery of the spiked metals resulted from complete (portlandite and brucite) or partial (gibbsite) dissolution of the host mineral. It could also be interpreted as reflecting the desorption of previously sorbed metals, which is likely to be the case for the high (60%) Cu recovery from hematite, where no significant dissolution of the host mineral was noted. In a similar study, Lo and Yang (1998) spiked discrete minerals (calcium carbonate, Fe and Mn oxides, bentonite, and kaolin), and humic acid with Pb, Cu, and Zn by thoroughly shaking them in solutions of the metals. The spiked components were mixed at different rates and diluted with fine sand to form soils. After aging for a week, the soils were extracted using the Tessier et al.’s (1979) SSE procedure (excluding the first step). As a result of the high metal recoveries obtained in the carbonate extraction step, the authors suggested that metals prefer to associate with the carbonate fraction rather than other fractions. However, possible contributions from specifically sorbed metals were not evaluated. Hence, specifically sorbed and carbonate-bound metals should be considered as two distinct entities. This is even more important in the case when amorphous free oxides constitute a major component of the specific adsorption capacity of the system. In such a case, occasional low redox potentials can affect free oxides solubility, rendering it a source rather than a sink for metals. Operationally identifying the metal in this step as carbonate-bound only is misleading. Ma and Uren (1998) solved this ambiguity by introducing an additional step (NH4 OAc + EDTA at pH 8.3) prior to extracting the carbonate-bound component (Table 1). Inasmuch as metals recovered in this step may not be referred to as a specific chemical component, they could


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be assigned functional attributes, such as mobility and bioavailability, analogous to the DTPA extraction for available trace elements (Lindsay & Norvell, 1978).


The Organic Component Metals associated with organic matter are extracted by destroying the organic material with hydrogen peroxide (H2 O2 ) or sodium perchlorate (NaClO4 ), or by dispersing the organic matter with complexing agents (e.g., DTPA, Na4 P2 O7 , EDTA; Beckett, 1989). Organic matter is also hydrolyzed and dispersed in alkaline media. In a recent thorough review, Mikutta et al. (2005) evaluated methods used to remove organic matter. Extraction with potassium, or sodium pyrophosphate (K4 , or Na4 P2 O7 ), and sodium hydroxide (NaOH; both for 16–24 hr at room temperature) or NaClO4 (for 30 min at 90–95◦ C) is performed at pH 8–10, whereas extraction with H2 O2 is done at pH 2 (for 3 hr at 80◦ C). Na4 P2 O7 and H2 O2 are less preferred extractants because they allow only partial dissolution of the organic material and are not selective enough. Extraction with H2 O2 at pH 2 or Na4 P2 O7 at pH 10 also dissolves manganese oxyhydroxides, amorphous iron oxyhydroxides, and sulfides (Beckett, 1989; Miller et al. 1986b; Shuman, 1983, 1991). Lavkulich and Wiens (1970) showed NaClO to be more effective and selective (i.e., destroying more organic matter while attacking less soil oxides) than H2 O2 as extraction for soil organic matter. Similarly, Shuman (1983) found that Na4 P2 O7 (at 0.1 M, pH 10) extraction dissolved less organic material than did NaClO4 . He also found that NaClO4 was much more selective than either H2 O2 or Na4 P2 O7 , which more readily dissolved portions of the Mn and iron oxides. Shuman (1982) found that Na4 P2 O7 (0.1 M, pH 10) extracted similar amounts of Fe, as did Tamm’s reagent (used to extract amorphous iron oxides; see subsequent section on oxyhydroxide). Hence, Shuman (1982, 1983) suggested that Na4 P2 O7 be used to characterize metals in the organic matter fraction only if preceded by extraction of the oxides. However, McKeague et al. (1971), working on amorphous and crystalline iron oxides, showed that 0.1 M Na4 P2 O7 dissolved a negligible amount of the minerals. They did find this extraction very efficient in solubilizing Fe–fulvic acid complexes (especially at Fe–C atomic ratio < 0.2). Hence, the fate of easily reducible oxides under the extraction conditions for organic matter must be taken into consideration when designing a SSE procedure. Indeed, the removal of the organic fraction with Na4 P2 O7 (Fitamo et al., 2007) or H2 O2 (Tessier et al., 1979) is often followed by the extraction of the free oxides to prevent their dissolution and hence their contribution to the organic component (Table 1). In protocols that manganese oxyhydroxides and amorphous iron oxides are characterized separately, the organic fraction is extracted after the manganese oxyhydroxides and before the amorphous iron oxides.


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Krishnamurti and Naidu (2000, 2002) included an extraction with 0.1 M Na4 P2 O7 at pH 10, which they named metal-organic complex-bound. This extraction was performed after removing the exchangeable metals (in 1 M NH4 NO3 at pH 7) and the carbonates (1 M NaOAc pH 5), and prior to extraction of the easily reducible and amorphous free oxides (Table 1). They further treated the Na4 P2 O7 extract with acid (pH 1) in order to separate between humic-acid-bound metals and fulvic-acid-bound metals by precipitating the former. The majority of the organic bound Cu, Zn, and Cd were associated with the fulvic acid component (71%, 78%, and 73%, respectively; Krishnamurti & Naidu, 2000, 2002). Inasmuch as this pH dependent of humic material is a common procedure for separating the two organic acid groups (Stevenson, 1994), the employed conditions (i.e., pH 1) are likely to release metals bound to humic acids, leading to an overestimation of the fulvic-acidbound metal component. The use of Na4 P2 O7 in the sequence before the extraction of the easily reducible component, and the acidification of the Na4 P2 O7 -extract may explain why the Cu and Zn fulvic acid components were the second largest fraction (after the residual fraction) found in these soils, and why the metal content in the easily reducible and the amorphous components were among the lowest. Another problem with this step, regardless of the method used, is the possible resorption and precipitation of metals after their release. Organic ligands (e.g., EDTA, DTPA) cannot be used in an oxidizing environment, and their addition at the end of the alkaline extraction caused dissolution of oxyhydroxides (Shuman, 1983). To address this problem, complementary subsequent extractions were added to this step at the end of the reaction, including the addition of 1 M MgCl2 (Shuman, 1979), 0.01 M HNO3 (Han & Banin, 1996), NH4 OAc (Quevauviller 1994, 1998a; Tessier et al., 1979), 1 M NH4 NO3 (Han et al., 2001), 2 M NH4 NO3 in 20% HNO3 (Krishnamurti & Naidu, 2000, 2002), or 1 M HOAc (Nyamangara, 1998). In some studies an additional substep follows the extraction of the organic sample. Sloan et al. (1997) and Burton et al. (2005) extracted the organic matter residue with NH4 OAc (3.2 M); Ma and Uren (1995, 1998), who extracted the organic component after the carbonates, repeated the carbonate extraction step (i.e., 0.5 M NaOAc, pH 4.74) after the organic component step. The extractable metals in these substeps were added to the organic fraction.

Oxyhydroxides Oxyhydroxides are separated into two or three main components in SSE protocols: manganese oxyhydroxides, amorphous iron oxyhydroxides, and well-crystallized iron oxyhydroxides (Table 1). Often, the first two are extracted together, and sometimes all three are extracted in one step. In some SSEs, the oxides are included in the residual component.


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Reducing agents, ligands, and acids are used to dissolve oxyhydroxides (Stumm & Furrer, 1987). A wide selection of agents, concentrations and extraction conditions (mainly time, pH, and temperature) are used to separate between the three oxyhydroxide types. As the reaction conditions become harsher, the dissolution encompasses more components. The ability to specifically and selectively dissolve these oxyhydroxide entities is based on controlling the reactions kinetics by specific sets of conditions (e.g., temperature, reaction time, concentrations, solid–solution ratio, irradiation). Hence, these set of conditions needs to be meticulously followed and reported.


MANGANESE OXYHYDROXIDES Mn-oxyhydroxides were shown to be specifically and selectively extracted in 0.1 M NH2 OH-HCl, (pH 2, at room temperature for 30 min; Chao, 1972; Shuman, 1982, Silveira et al., 2006). This step is performed after the dissolution of the carbonates and before the extraction of the organic fraction (Table 1). Ma and Uren (1995, 1998) extracted the manganese oxyhydroxides with 0.2% hydroxyquinone in 1 M NH4 OAc at a neutral pH. This extraction preceded the carbonate extraction because of a concern that the pH used to extract carbonates (4.75 in their protocol) dissolves manganese oxyhydroxides as well.

AMORPHOUS IRON OXYHYDROXIDES Amorphous Fe-oxyhydroxides are selectively extracted by Tamm’s reagent (also known as acid ammonium oxalate [AAO]), which consists of 0.175 M ammonium oxalate (NH4 Ox), 0.1 M oxalic acid (HOx) at pH 3.25. The dissolution is performed in the dark by shaking for 2–4 hours at room temperature (Chao & Zhou, 1983; Endredy, 1963; Fine & Singer, 1989; Loeppert & Inskeep, 1996; Schwertmann, 1973, 1991; Shuman, 1991). Chao and Zhou (1983) found that extraction with 0.25 M NH2 OH-HCl in 0.25 M HCl at 50◦ C for 30 min corresponded well with the Tamm’s reagent in terms of selectivity and specificity. Because the former is much less time consuming, Chao and Zhou (1983) preferred it over Tamm’s reagent. However, Ross et al. (1985) found no correlation between the two extractions in the 27 soils that they examined; the Chao and Zhou (1983) procedure consistently had low recoveries. They did obtain similar results for the two methods when the Chao and Zhou (1983) reagent reacted with the soils for 16 hr at room temperature instead of 30 min at 50◦ C. Wang et al. (1987) also showed AAO extraction to be more effective than NH2 OH-HCl in extracting amorphous iron and aluminum from spodosols. Similar results were obtained by Benitez and Dubois (1999) and Davidson et al. (2004) for Mn and iron oxyhydroxides. However, they found that recovery of heavy metals, especially Cd and Pb, were lower with AAO. Elliott and Herzig (1999) demonstrated a higher efficiency in the removal of Pb from contaminated soil under acidifying conditions in the absence, rather than in the presence of 1 M HOx. Less conclusive results


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were obtained for Zn extraction. The lower recovery of heavy metals with AAO was attributed to re-precipitation of insoluble metal-oxalate complexes under acidic conditions (Elliott & Herzig, 1999). A major difficulty concerning extraction of Mn oxyhydroxides and amorphous iron oxyhydroxides is that the acidic conditions of these reactions overlapped those employed for the extraction of the organic component. This brought about a reordering of the SSE sequences. Hence, in studies in which Mn oxyhydroxides and amorphous iron oxyhydroxides were extracted in one step, the organic component extraction was extracted in a subsequent step (e.g., Gupta & Chen, 1975; Sloan et al., 1997; Table 1). Otherwise, Mn oxyhydroxides were extracted separately before the extraction of the organic component, which was then followed by extraction of the amorphous iron oxyhydroxides (e.g., Berti & Jacobs, 1996; Miller et al., 1986b; Table 1).

WELL-CRYSTALLIZED IRON OXIDES Well-crystallized Fe oxides are often extracted using sodium dithionite (Na2 S2 O4 ) in a citrate-bicarbonate buffer (CBD; Mehra & Jackson, 1960). This is the preferred extraction procedure for the specific and selective dissolution of iron oxyhydroxides, excluding magnetite (Arshad et al., 1972; Borggard, 1988; Chao & Zhou, 1983; Fine & Singer, 1989; McKeague & Day, 1966; Walker, 1983). However, because of metal impurities (mostly zinc; Shuman, 1982), and possible precipitation of extracted metals as sulfides during this procedure, this reagent is rarely used in SSE to study heavy metal distribution in soils and other media (e.g., used by Gibbs, 1973; Mandal & Mandal, 1986). Instead, the Tamm’s reagent is often used under UV irradiation (at 80–100◦ C for 3 hr; Beckett, 1989; Endredy, 1963) or extraction is done with 0.04 M NH2 OH-HCl in 25% HOAc at pH 2 and 96–100◦ C (Beckett, 1989; Tessier et al., 1979). Shuman (1982) preferred using solution of 0.2 M NH4 Ox, 0.2 M HOx, 0.1 M Ascorbic Acid (pH 3) at 97◦ C because of the impurities and precipitation problems of the CBD extraction on the one hand, and the lack of standardization of the UV irradiation, on the other hand. Silveira et al. (2006) solubilized crystalline iron oxides in tropical soils by shaking them in 6 M HCl for 24 hr at room temperature (Table 1). This was preferred over the NH4 Ox-ascorbic acid at pH 3 method, which had to be repeated five times for comparative results. Noteworthy, the crystalline iron oxides extraction efficiency in 6 M HCl strongly decreases with an increasing degree of Al substitution, which occurs extensively in pedogenic goethites (Schwertmann, 1991). This should be a concern in the more weathered soils of the tropics. It should be noted, however, that CBD seems to dissolve secondary (mostly pedogenic) iron oxyhydroxides more efficiently then Tamm’s reagent or NH2 OH-HCl, HOAc at 96◦ C (Rozan et al., 1997). This was further confirmed in La Force and Fendorf’s (2000) comparative study, which showed that NH2 OH-HCl had low dissolution efficiency for well-crystallized



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FIGURE 5 Dissolution of synthetic iron oxyhydroxides by 0.04M NH2 OH-HCl in 25% HOAc for 6 hr as a function of temperature. = goethite; •, = ferrihydrite; = lepidocrocite (reprinted from Geochim, Cosmochin. Acta, 57, 4391–4404, Fortin et al., Characteristics of lacustrine diagenetic iron oxyhydroxides, 1993, with permission from Elsevier).

iron oxyhydroxides. Fortin et al. (1993) also showed low efficiency of 0.04 M NH2 OH-HCl in 25% HOAc at 96◦ C in dissolving synthetic goethite (