CHAPTER 2 INTRODUCTION

Transcription

1 CHAPTER 2 Phosphate Adsorption by Mixed and Reduced Iron Minerals Under Static Conditions INTRODUCTION I. Importance and cycling of phosphate Phosphorus (P) availability is an important factor in biogeochemical cycling and is often the limiting nutrient in freshwater and estuarine ecosystems (Reddy et al., 1999). It is a macronutrient required by plants and animals and is a critical component in DNA, RNA, ATP, and phospholipids (Schlesinger, 1991). Too little available P in agricultural systems can result in poor productivity and crop failure. Conversely, increased P input to fresh or saltwater systems via erosion can lead to large algal blooms and subsequent eutrophication of water bodies (Brady and Weil, 1999). Noe et al. (2001) have recently documented that even small P increases due to increased agricultural runoff in the Everglades appears to be altering species distribution and, subsequently, the ecology of the system. The ultimate source of P to the biosphere is weathering of rocks, and the amount required for biological production is much greater than the amount released by weathering. This deficit results in P being highly conserved and rapidly cycled within the biosphere (Filippelli and Souch, 1999; Berner and Berner, 1996). Vegetation, periphyton, plankton, and microbes all act as biotic phosphorus sinks. However, biotic cycling is not without P loss to abiotic mechanisms, including adsorption by soils and sediments, mineral precipitation, and sedimentation. Minerals controlling P in soils and sediments include apatite, hydroxyapatite, fluorapatite, octocalcium phosphate, strengite, vivianite, variscite, and wavellite (Reddy et al., 1999). Phosphorus adsorption by soils and lake sediments has often been correlated to amorphous iron, aluminum, and manganese oxides and hydroxides, as well as calcium and magnesium (Brady and Weil, 1999; Reddy et al., 1999; Torrent, 1997; Golterman 1995; Olila and Reddy, 1995; Blanchar and Frazier, 1991; Williams et al., 1971). The relationship between P and ferric oxides is usually more apparent in acidic to neutral soils, while in alkaline soils P is often associated with calcium or magnesium (Reddy et al., 1999). 33

2 II. Phosphorus retention processes The prominent correlation between phosphorus retention and iron has led to numerous examinations of P adsorption capacities and mechanisms onto iron oxides and hydroxides in aerobic environments. The most commonly studied ferric minerals have been hematite and goethite since they are the most abundant, although not necessarily the most reactive, iron minerals found in soils (Torrent et al., 1994; Sparks, 1995). The mechanism of phosphate adsorption onto ferric oxides is dominated by a ligand exchange reaction in which two singly coordinated hydroxyls or water molecules are replaced by a single phosphate anion resulting in the formation of a bidentate, binuclear complex (Reddy et al., 1999; Torrent, 1997; Colombo et al., 1994; Parfitt and Russell, 1977; Parfitt et al., 1975). Parfitt et al. (1975) used infrared (IR) spectroscopy to verify the formation of binuclear complexes on lepidocrocite, goethite, hematite, ferric hydroxide gels, and β-ferric hydroxide. Arai and Sparks (2001) used attenuated total reflectance Fourier transform infrared (ATR-FTIR) spectroscopy to show that the P adsorption mechanism on ferrihydrite was a non-protonated, binuclear complex at ph greater than 7.5. Binuclear complexes are quite strong and result in slow exchange rates and an apparent irreversibility of phosphorus adsorption. The strength of these complexes lead to long-term phosphorus storage in soils, sediments, and wetlands (Reddy et al., 1999; Parfitt et al., 1975). However, formation of mononuclear complexes cannot be overlooked. Torrent (1997) stated that P formed mono- and binuclear complexes on goethite at low ph and binuclear complexes were only predominant near neutral or higher ph. Nonetheless, it is clear that phosphorus forms very stable surface complexes on iron (hydr)oxide solids. Two stages of phosphate sorption transpire. Rapid adsorption takes place during the first few minutes of the reaction and is followed by a slower stage, which may be measurable for days to weeks, attributed to diffusion into the solid phase (Reddy et al., 1999; Strauss et al., 1997; McLaughlin et al., 1977). Torrent (1997) summarized several other proposed theories explaining the slow adsorption step. These include the initial formation of mononuclear complexes and rearrangement to form binuclear complexes, replacement of silicate in the sorbent, competing anions on the surface, and surface precipitation processes. Slow adsorption increases with specific surface area, micro- and 34

3 meso-porosity, and ferrihydrite impurities (Strauss et al., 1997; Torrent 1997; Torrent et al., 1994). This slow sorption has also been associated with irreversible phosphate retention (Madrid and DeArambarri, 1985). Phosphate adsorption capacity generally averages 2.5 µmols P/m 2 for most iron oxides, ranges from 1.5 to 3.5 µmols/m 2 (Reddy et al., 1999; Torrent 1997; Enyard, 1994; Torrent et al., 1990; Borggaard, 1983). Surface area and crystallinity are the best indicators of phosphorus retention capacities (Strauss et al., 1997; Colombo et al., 1994; Enyard, 1993; Borggaard, 1983; McLaughlin et al., 1981). Phosphate adsorption on synthetic iron oxides has been described by the Langmuir and Freundlich isotherms (Miltenburg and Golterman, 1998; Torrent, 1997; Golterman, 1995; Gerke, 1993; McLaughlin et al., 1977). Colombo et al. (1994) found adsorption on hematite ranged from 0.31 to 2.27 µmols P/m 2 and was described by the Freundlich isotherm. This range was attributed to variations in crystal morphology and size, and it is supported by the work of Torrent et al. (1994), where adsorption capacities of natural hematite were reported between 0.8 and 4.1 µmols P/m 2. Sorption onto hematite was more variable than onto goethite. The rate and adsorption capacity of phosphate on goethite is highly dependant on crystallinity (Strauss et al., 1997). Crystalline goethite with a surface area of 18 m 2 /g had an adsorption capacity near 2.5 µmols P/m 2, and adsorption was complete within a day. Several other less crystalline goethite samples reacted as long as 3 weeks and exceeded the 2.5 µmols P/m 2 adsorption capacity. In general, it appears that phosphorus sorption by ferric hydroxide gel and goethite decreases with increasing ph (Geelhoed, 1997; Willett and Cunningham, 1983; Ryden et al., 1977a; McLaughlin et al., 1977; Hingston et al., 1972). Strauss et al. (1997) found P adsorption is greatest at ph near 2. Shang et al. (1992) suggested ph has an appreciable influence on reaction rate in the beginning of sorption reactions and noted an overall decrease in phosphorus adsorption capacity with ph from ph 4 to 6. Strauss et al. (1997) also observed P sorption onto goethite occurred most rapidly at ph 2 and that ph was particularly important during the initial fast reaction, with high ph slowing the reaction kinetics. Shang et al. (1992) examined the kinetics and overall adsorption of phosphate by amorphous iron precipitates and also found that as ph increased sorption 35

4 decreased with the most prominent difference in the fast adsorption step, defined as 0.5 h. The authors also found that total phosphate sorption decreased with ph and attributed this occurrence to the presence of fewer fully protonated (i.e, water) leaving groups, which are preferred over hydroxyl groups. Torrent (1997) stated that phosphate formed monoand binuclear complexes on goethite at low ph and binuclear complexes were only predominant near neutral or higher ph. Arai and Sparks (2001) found phosphate sorption on ferrihydrite decreased with ph from 3.5 to 9. Arai and Sparks (2001) also noted that at ph greater than 7.5, phosphate formed non-protonated, bidentate, binuclear complexes on ferrihydrite, while from ph 4 to 6 their evidence suggested the presence of protonated complexes. Understanding how the sorption characteristics of a given oxide change with ph is important when considering phosphate cycling. Oxides with well defined, sharp adsorption envelopes are likely to be more sensitive to ph fluctuations and are less likely to strongly retain phosphate under changing ph conditions. Ionic strength may also have an important influence on phosphate retention. Ryden et al. (1977b) and McLaughlin et al. (1977) found P adsorption onto ferric hydroxide gel to increase with ionic strength. Arai and Sparks (2001) found that from ph 4 to 7.5, P adsorption on ferrihydrite was unaffected by changes in ionic strength; however, at ph greater than 7.5, adsorption increased with ionic strength. Work by Celi et al. (2000) on goethite has shown increased ionic strength results in increased phosphorus sorption as well. However, Geelhoed et al. (1997) found that increasing the ionic strength at low ph resulted in lower P adsorption but at higher ph an increase in sorption was observed. III. Needs and Research Objectives Many studies have discounted iron oxides as playing a significant role in P retention under anoxic conditions, suggesting P will be released to solution upon the reductive dissolution of the iron (III) oxides or phosphates. However, Patrick and Khalid (1974) found the adsorption capacity for iron-rich sediments increases under reducing conditions. They postulated that reductive dissolution of iron oxides and hydroxides and subsequent precipitation of ferrous minerals with high surface areas increased P adsorption capacity in an aerobic environment, but that the P would be less tightly held 36

5 than in the aerobic environment. Thus, there are conflicting sentiments regarding P retention in anaerobic environments. Accordingly, the goal of this study is to elucidate the binding capacity of phases common to anaerobic environments for P retention. Specifically, we investigate P sorption on a number of biogenic ferrous-bearing minerals that may be formed by microbial reduction of ferric oxides and hydroxides in anaerobic environments. Identified ferrous phases include siderite, vivianite, magnetite, and green rust (Fredrickson et al., 1998; Zachara et al., 1998). Characterizing P retention by these minerals is important since microbially and chemically induced mineral transformations within soils, sediments, and riparian zones can have a substantial influence on the mobility of P along with other nutrients, and contaminants (Loyaux-Lawniczak et al., 2000; Erbs et al., 1999; Myneni et al., 1997; and Hansen et al., 1996). Riparian zones are often important in determining the behavior of contaminants in freshwater systems, since these areas can be pathways between surface run-off and rivers. Changing oxidation reduction conditions in riparian zones will dictate the formation of specific iron minerals, partially through biogenic processes. Results from this study will help determine how phosphate interacts with iron phases in suboxic environments. More importantly, this study will help to elucidate phosphorus dynamics within riparian zones, thus improving our understanding of the fate and transport of phosphate in the freshwater environment. Knowledge of the extent of P retention by mixed and reduced iron minerals will also have implications in nutrient management from an agricultural standpoint, particularly with respect to fields subjected to periodic flooding. Alteration between aerobic and anaerobic conditions could lead to formation of the mixed and reduced iron minerals and these phases will need to be considered when making decisions about nutrient management and when evaluating P dynamics in a system. The specific objectives of this study are (1) to determine P sorption capacities of siderite, magnetite, and sulfate green rust and compare retention to that of lepidocrocite, goethite-coated sand, ferrihydrite-coated sand, and ferrihydrite slurry, (2) to determine the influence of ph and ionic strength on retention capacity, and (3) to estimate P sorption within anaerobic and aerobic soils. Sorption isotherms were determined using batch experiments allowed to react for 48 hours. Initial P concentrations from 10 to

6 µm were evaluated. Ionic strength and ph effects were investigated using 100 µm P solutions with ionic strength of 0.05 M to 0.1 M and ph from 4.5 to 9. 38

7 MATERIALS AND METHODS Batch studies were conducted using several mixed and reduced iron minerals as sorbents. Natural and synthetic samples of magnetite and siderite were studied in addition to green rust, ferrihydrite, ferrihydrite coated sand, goethite coated sand, and lepidocrocite. I. Mineral preparation Naturally occurring siderite was purchased from Ward s Scientific (46E7351, Antigonish County, Nova Scotia) and naturally occurring magnetite was obtained from the Stanford University Research Mineral Collection. Synthetic magnetite was purchased from Aldrich Chemical ( ). The natural (of geologic origin) magnetite and siderite samples were crushed and passed through a 250 µm sieve. Both magnetite samples were treated with 1 M HCl for 1 h to remove oxidized surface constituents, rinsed three times with anoxic, deionized water and dried using vacuum filtration. The siderite was treated using 0.05 M HCl for 1 h, rinsed, and dried. All mixed or reduced iron minerals were stored in an anaerobic glovebox (Coy Products, Ann Arbor, MI) having O 2(g) levels less than 5 ppm. Synthetic siderite was made inside an anaerobic glovebox using the method described by Hallbeck et al. (1993). Anoxic water was used to prepare 0.1 M Fe(NH 4 ) 2 (SO 4 ) 2 6H 2 O and 0.1 M Na 2 CO 3 solutions. Water was made anoxic by boiling and bubbling with oxygen-free N 2 for 1 h prior to making solutions. A cloudy white precipitate formed immediately upon the addition of 500 ml of the Na 2 CO 3 solution to 500 ml of the Fe(NH 4 ) 2 (SO 4 ) 2. The solid was allowed to settle, rinsed three times with anoxic deionized water, and dried using vacuum filtration followed by desiccation. Ferrihydrite (2-line) was made using the method described by Cornell and Schwertmann (1996). A ferric chloride solution (2.3 L of M FeCl 3 6H 2 O, ph near 2) was titrated with 0.2 M NaOH to ph 7. The solution was then stirred and allowed to stabilize for several hours with ph maintained by base addition. Once the solution ph was stable, the precipitate was allowed to settle and rinsed three times with deionized water. Ferrihydrite coated sand was made by drying the slurry onto 660 g of Iota 6 39

8 Quartz Sand (Umimin Corporation). The sand was dried at room temperature, rinsed several times with deionized water, and dried again. Goethite was synthesized using a method described by Cornell and Schwertmann (1996). A 1 M NaHCO 3 (110 ml) was added to 1L of 0.5 M FeCl 2 4H 2 O (made with deoxygenated water) while the latter was stirred and bubbled with N 2. Upon addition of the NaHCO 3 the flow of N 2 was replaced with air at a rate of ml/minute. The solution was stirred and bubbled for 48 h. When oxidation was complete, solids were allowed to settle and the supernatant decanted. The remaining slurry was then washed by suspension in a 0.1 mm NaCl solution followed by centrifugation. This rinse procedure was repeated a total of five times and after the last rinse the slurry was sonicated for 1 h. Goethite-coated sand was made by drying the slurry onto 278 g of Iota 6 Quartz Sand (Umimin Corporation). The sand was dried at room temperature, rinsed several times with deionized water, and dried again. Lepidocrocite was synthesized following a method described by Cornell and Schwertmann (1996). A 0.6 M FeCl 2 solution was titrated to ph 6.5 using 1 M NaOH. Air was then introduced to the solution by bubbling at a rate of ~250 ml/min. The ph was maintained between 6.3 and 6.5 using 1 M NaOH and the process terminated when no additional base was necessary to maintain ph. Sulfate green rust was precipitated using a modification of the method described in Schwertmann and Fetcher (1994). A 0.5 M ferrous sulfate solution (ph near 3.7) was prepared using anoxic deionized water. The ferrous sulfate solution was kept in an anoxic reaction vessel and the ph was brought to 7.1 using 4 N NaOH. Air was then introduced at 60 ml/min while additions of 4 N NaOH or 1 N HCl were added to maintain the ph. The reaction was complete when the E h stabilized. II. Mineral characterization X-ray diffraction was carried out using a Rigaku Geigerflex diffractometer (Model CN2029) with Cu Kα radiation (35 kv, 15 ma). Samples were prepared for analysis by pressing the powdered mineral material into a 0.5 mm depression on a Rigaku mono-crystalline silica XRD slide. The synthetic siderite was easily oxidized and thus the slide needed to be prepared in the glovebox. In order to prevent oxidation during 40

9 analysis, a small amount of the solid was mixed with glycerol to form a thick paste, which was then packed into the depressed slide. The sample was analyzed immediately upon removal from the glovebox. Specific surface area was determined by N 2 adsorption using the BET equation (Brunauer et al., 1938). The adsorption curves were obtained using a Coulter SA 3100 surface area analyzer. All minerals, except synthetic siderite and sulfate green rust, were degassed by heating overnight at 85 C. Contact with the atmosphere was avoided by maintaining positive pressure (helium) on the sample containers. Excess water was removed from the synthetic siderite and green rust by dessication under a vacuum for 1 week. III. Phosphate adsorption experiments Solutions used to define the adsorption isotherms ranged from 10 µm to 1500 µm PO 4 and were made using Na 2 HPO 4. Ionic strength was adjusted to 0.1 M using NaCl. The ph was adjusted to 7 using anoxic 0.05 M HCl or 0.05 M NaOH. The ph effects were evaluated using 100 µm PO 3-4 solutions adjusted to ph 4, 5.5, 7, 8, and 9. Ionic strength influences were investigated using 100 µm PO 3-4 solutions at ph 7 with the ionic strength adjusted between 0.1 and M. All solutions used with mixed or reduced iron minerals were made under anaerobic conditions. All experiments using mixed or reduced iron minerals were prepared in an anaerobic atmosphere. Acid-cleaned glassware was allowed to equilibrate in the glovebox at least 12 h before use. Approximately 0.1 g mineral was weighed into 125 ml serum vials, followed by the addition of 100 ml of solution. The supply of green rust was limited, so about 0.02 g mineral was weighed into a 25 ml serum vial and 15 ml of phosphate solution added. The vials were sealed, removed from the glovebox, and shaken using an orbital shaker (120 rpm) for 48 h. All samples were run in triplicate, with the exception of the sulfate green rust, which were run in duplicate. Upon completion, 8 ml of solution were passed through a 0.45µm filter and acidified using two drops of HNO 3. Starting and ending solutions were analyzed using inductively coupled plasma-optical emission spectroscopy (ICP-OES). The amount of phosphate adsorbed was determined from the difference between the initial and final concentrations. 41

10 IV. Release of sulfate from green rust Sulfate release from green rust with increasing ionic strength was evaluated by allowing the mineral to react for 48 h with deionized water and anoxic NaCl solutions ranging in ionic strength from to 0.1 M. Approximately 0.02 g green rust was weighed into a 25 ml serum vial, 15 ml solution added, and the sample placed on an orbital shaker (120 rpm). A solution containing only 100 µm NaH 2 PO 4 was also assessed using the same procedure. 42

11 RESULTS I. Adsorption capacities and isotherms Phosphate adsorption capacities for several minerals were higher than expected. Adsorption capacities were determined by measuring the amount of phosphate in solution before and after reaction with a known quantity of mineral. The amount of phosphate lost from solution was assumed to be adsorbed by the mineral phase. Experimental adsorption capacities ranged from 4.73 to µmols PO 4 /g mineral (Table 2.1). Goethite coated sand exhibited the lowest adsorption capacity and ferrihydrite the greatest capacity on a mass basis. Mineral surface areas ranged from 0.13 to approximately 200 m 2 /g. On a surface area basis, sulfate green rust exhibits the greatest capacity for phosphate ( µmols PO 4 /m 2 ) and lepidocrocite the lowest (1.65 µmols PO 4 /m 2 ). Surface area values and normalized sorption maxima are summarized in Table 2.2. Adsorption curves for geologic magnetite, green rust, and geologic siderite (Figures 1a-c) were best described by the Freundlich equation (R 2 = 0.88, 0.99, and 0.97, respectively). The linear form of the Freundlich equation is log (x/m) = (1/n) log C + log K [2.1] where K is the distribution coefficient, C is the equilibrium concentration, x/m is mass phosphate adsorbed/mass mineral, and n is a correction factor. The calculated distribution coefficients (K) were (magnetite), (green rust), and (siderite). Correction factors (n) for the three minerals were 1.38, 1.35, and 1.47 respectively. The Langmuir equation best describes the adsorption curves for ferrihydrite, ferrihydrite-coated sand, goethite-coated sand, and lepidocrocite (Figure 2.2), as well as synthetic siderite and synthetic magnetite (Figure 2.3). The calculated monolayer capacities and binding coefficient values for these equations are listed in Table

12 Table 2.1. Maximum observed phosphate adsorption on a mass basis of the mineral adsorbent (µmols PO 4 /g mineral). Maximum observed Mineral adsorption (µmols/g) SD* Goethite coated sand Natural magnetite Ferrihydrite coated sand Natural siderite Synthetic magnetite Lepidocrocite Sulfate green rust Synthetic siderite Ferrihydrite *SD = standard deviation. Table 2.2. Surface area (SA) and SA normalized adsorption maxima for phosphate on various iron oxides. Mineral Surface Area (m 2 /g) Maximum adsorption normalized to SA (µmols PO 4 /m 2 ) Sulfate green rust Natural magnetite Natural siderite Synthetic magnetite Synthetic siderite Ferrihydrite-coated sand Ferrihydrite slurry 200 (estimated) 5.19 Goethite sand Lepidocrocite Table 2.3. Summary of adsorption capacities (b-values) and K L -values for minerals conforming to the Langmuir isotherm. Mineral Monolayer capacity (µmols /g) Binding coefficient (K L ) Ferrihydrite slurry Synthetic siderite Lepidocrocite Synthetic magnetite Ferrihydrite-coated sand Goethite sand

13 II. Influence of ph and ionic strength Generally P sorption decreased with increasing ph, at least beyond a critical value (Figure 2.4). Goethite- and ferrihydrite-coated sand both exhibited linear decreases in adsorption. Natural siderite and natural and synthetic magnetite all exhibit a relatively well-defined adsorption envelope, with minimal P retention at ph 9. Lepidocrocite appears to behave similarly to both magnetite specimens and natural siderite; however, sorption was not examined at ph 9 and the envelope thus cannot be constrained. Ferrihydrite and sulfate green rust showed only slight decreases in adsorption with increased ph. Synthetic siderite noticeably deviates from the general trend. Phosphate adsorption reaches a maximum at ph 7 and there was little difference in sorption at ph 4 and 9. Overall, changes in ionic strength had a negligible influence on P adsorption within the ph range examined. The only appreciable change in sorption observed with ionic strength was on sulfate green rust (Table 2.4). III. Release of sulfate from green rust Sulfate release from the interlayer of the green rust is likely to be an important consideration when attempting to evaluate P adsorption. Hansen and Poulsen (1999) found that approximately 50% of the interlayer sulfate was replaced by phosphate, ultimately leading to the formation of vivianite. The release of sulfate from green rust increased from an average of 245 µmols SO 4 /g in deionized water to 632 µmols SO 4 /g in 0.1 M NaCl solution (Figure 2.5). The amount of sulfate released into the 0.1 M ionic strength phosphate solutions used to establish sorption isotherms ranged from 437 to 677 µmols SO 4 /g (Figure 2.6). The average sulfate released into 100µM NaH 2 PO 4 solution with no NaCl was 306 µmols SO 4 /g and release into 100µM NaH 2 PO 4 with ionic strength adjusted to 0.1 M using NaCl was 558 µmols SO 4 /g (Figure 2.6). It should also be noted that phosphate adsorption was greater, 95.8 µmols PO 4 /g, when the solution contained no NaCl. Adsorption averaged 66.3 µmols PO 4 /g when the ionic strength was adjusted to 0.1 M using NaCl. 45

14 Table 2.4. Average adsorption (µmols PO 4 /g) at high and low ionic strength. Comparisons were done using an initial 100 µm NaH 2 PO 4 solution at ph 7. Mineral 0.05 M SD* 0.1 M SD* Green rust Natural magnetite Ferrihydrite Synthetic magnetite Lepidocrocite Goethite coated sand Synthetic siderite Ferrihydrite coated sand Natural siderite *SD = standard deviation. 46

15 a log (x/m) log C x/m (mg PO4 / gram mineral) log (x/m) log C b c log (x/m) log C C (equilibrium concentration µm PO 4 ) Figure 2.1. Phosphate adsorption curves for minerals fit to a Freundlich isotherm (linear form is inset): (a) geologic magnetite (R 2 = 0.89), (b) sulfate green rust (R 2 = 0.99), and (c) geologic siderite (R 2 = 0.97). In the case of sulfate green rust data points obscure the error bars. Observed adsorption maxima were: µmols PO 4 /g magnetite, µmols PO 4 /g GR, and µmols PO 4 /g siderite. 47

16 100 a b x/m (mg PO4 / gram mineral) C/(x/m) C c C/(x/m) C d C/(x/m) C C/(x/m) C C (equilibrium concentration µm PO 4 3- ) Figure 2.2. Adsorption isotherms of ferric (hydr)oxides all fit to the Langmuir isotherm (linear from in inset). (a) ferrihydrite (R 2 = 0.99), (b) ferrihydrite-coated sand (R 2 = 0.99), (c) goethite-coated sand (R 2 = 0.99), (d) lepidocrocite (R 2 = 0.99). Error bars in a are obscured by data points. The tabulated K l and b-values are listed in Table 3. Observed maxima were: µmols PO 4 /g ferrihydrite, µmols PO 4 /g ferrihydrite-coated sand, 4.73 µmols PO 4 /g goethite-coated sand, and µmols PO 4 /g lepidocrocite. 48

17 100 a Adsorbed P Concentration, x/m (mg PO4/g solid) C/(x/m) C b C/(x/m) C Equilibrium PO 4 concentration, C (µm) Figure 2.3. Phosphate adsorption curve fit to a Langmuir isotherm for (a) synthetic siderite (linear form in inset, R 2 = 0.96) and (b) synthetic magnetite (linear form in inset, R 2 = 0. 98). The tabulated K l and b-values are listed in Table 3. Observed adsorption maxima were µmols PO 4 /g siderite and µmols PO 4 /g magnetite. 49

18 14 12 Ferrihydrite slurry 10 Synthetic siderite Adsorbed P Concentration, x/m (mg PO4/g solid) Green rust Synthetic magnetite Ferrihydrite sand Goethite sand Geologic Siderite ph Geologic magnetite Figure 2.4. Influence of ph on phosphate adsorption to various iron oxides. Comparisons were done using an initial 100 µm NaH 2 PO 4 solution with I = 0.1 M. 50

19 70 mg SO 4 Released (mg SO 4 /g GR) No phosphate NaCl Concentration (M) 100 µm PO 4 (no NaCl) Figure 2.5: Sulfate released from green rust with increasing ionic strength (as NaCl). 51

20 Concentration (mg/g GR) PO 4 sorbed SO 4 released 0 PO 4 sorbed at 100 µ PO 4, no NaCl SO 4 released at 100 µ PO 4, no NaCl Equilibrium Phosphate Concentration (µm) Figure 2.6. Phosphate adsorption and sulfate release with increasing phosphate concentration. Also shown is phosphate sorbed and sulfate released into 100 µm phosphate solution when no NaCl is present. 52

21 DISCUSSION Sorption capacities for the ferric iron minerals examined in this study ranged from 1.65 to 6.31 µmols PO 4 /m 2. Ferrihydrite coated sand had the greatest capacity (6.31 µmols PO 4 /m 2 ), followed by ferrihydrite (5.19 µmols PO 4 /m 2 ), goethite-coated sand (1.98 µmols PO 4 /m 2 ), and lepidocrocite (1.65 µmols PO 4 /m 2 ). A review by Torrent (1997) listed the ranges of adsorption capacities for ferrihydrite ( µmols PO 4 /m 2 ), goethite ( µmols PO 4 /m 2 ), and lepidocrocite ( µmols PO 4 /m 2 ). The experimental results from this work are comparable with regard to lepidocrocite; however, the results for the remaining ferric phases are not as easily compared to the previously determined values. The adsorption capacity determined in this study for ferrihydrite is higher than those cited by Torrent (1997). This is likely the result of underestimating the surface area, which has been found to range from 100 to 700 m 2 /g (Cornell and Schwertmann, 1996). A value of 200 m 2 /g was assumed during the data analysis of this work. Based on the sorption capacity per weight of ferrihydrite in this study, a surface area of ~400 m 2 /g, well within the documented range, would have resulted in a value near that previously published. A batch sorption experiment conducted on uncoated silica sand indicated that the sand alone adsorbed no measurable phosphate; therefore, phosphate adsorption by the goethite- and ferrihydrite-coated sands is due entirely to the respective iron phase. This is also reflected in the results of the surface area analyses. Uncoated silica had a surface area of 0.14 m 2 /g, while ferrihydrite- and goethite-coated sand yielded values of 2.28 m 2 /g and 2.39 m 2 /g, respectively. Iron content of the sand (weight percent) was 0.35% for the ferrihydrite-coated sand and 0.42% for the goethite-coated sand. When comparing the sorption capacities of the ferrihydrite and ferrihydrite-coated sand normalized to iron (weight percent), sorption by the coated-sand is much greater (~2000 µmols PO 4 /g Fe as compared to ~4100 µmols PO 4 /g Fe). This may be the result of aggregation of the ferrihydrite slurry (no sand), which would decrease the number of sorption sites readily available. Phosphate sorption by all four ferric iron minerals was fit to a Langmuir isotherm. The initial slopes of the adsorption curves for lepidocrocite, ferrihydrite, and goethite- 53

22 coated sand were quite steep; however, the initial section of ferrihydrite-coated sand adsorption curve is more gradual. These observations are reflected in the values of the binding coefficients (K L ), which are for ferrihydrite and goethite coated sand and for lepidocrocite. The K L for ferrihydrite-coated sand was The steep adsorption curves and high binding coefficients indicate a high affinity for phosphate by lepidocrocite, ferrihydrite, and goethite-coated sand, while the ferrihydrite-coated sand has a lower affinity for phosphate on a mass basis. The reduced and mixed valent iron phases studied exhibited sorption capacities higher than the oxidized minerals when normalized to surface area, ranging from to µmols PO 4 /m 2. Synthetic magnetite and synthetic siderite reach adsorption maxima of and µmols PO 4 /m 2, respectively. The binding coefficients are similar to that of ferrihydrite-coated sand with the K L for siderite calculated at and magnetite at The slopes of the initial section of the adsorption curves reflect these lower binding affinities, which are similar to that of ferrihydrite-coated sand. The natural magnetite, natural siderite, and sulfate green rust conformed to the Freundlich isotherm and thus no adsorption maximum was observed. This is likely explained by surface precipitation and/or physical adsorption of phosphate in addition to chemisorption at the mineral surface. To evaluate whether surface precipitation occurred, samples of natural siderite, natural magnetite, and sulfate green rust were reacted with 1500 µm NaH 2 PO 4 (ph 7 and ionic strength adjusted 0.1 M using NaCl) for 48 h, dried, and analyzed using X-ray diffraction. The ending solution was analyzed for phosphate and dissolved iron; the chemical speciation program MINTEQ was used to determine saturation indices using parameters from Stumm and Morgan (1981) and Benjamin (2002) (Table 2.5). X-ray diffraction patterns did not show evidence of iron phosphate mineral phases; however, if surface precipitates were present as a small percentage of the sample they would not be detected. Vivianite was found to be oversaturated in solutions reacted with magnetite and siderite, so it is possible for precipitation to occur. The solution reacted with sulfate green rust was slightly undersaturated with respect to vivianite (Table 2.5); however, based on the work by Hansen and Poulsen (1999) conversion of green rust to vivianite is a likely phosphate sink in these systems. They reacted sulfate green rust with 50 mm Na 2 HPO 4 (ph 9.3) for 60 d. 54

23 Table 2.5. Concentration data and saturation indices for phosphate solutions after 2 d reaction with geologic magnetite, geologic siderite, and sulfate green rust. Mineral Fe (mm) P (mm) SI* Fe(OH) 2 SI* Vivianite SI* Wustite Magnetite A Magnetite B Green Rust A Green Rust B Siderite A Siderite B *SI= saturation index (log IAP/K sp, where IAP is the ion activity product for the respective phase) Vivianite was detected within 1 d l of phosphate addition and most of the vivianite precipitation occurred during the first 10 d of the reaction; however, continued precipitation was noted after 40 d. The first phase of precipitation was attributed to vivianite formation resulting from the reaction of Fe(II) in solution with the added phosphate. The later phase was believed to follow a period of phosphate reaction with Fe(II) within the green rust resulting in vivianite precipitation in the interlayer. They also found that the interlayer thickness decreased and that sulfate was released into solution as a result of exchange with phosphate. The authors also stated that even though phosphate does exchange with sulfate in the interlayer, the affinity for phosphate is not great enough to lead to complete substitution. Nevertheless, the continued reaction observed here may be attributed to phosphate reacting in the interlayer of green rust. Sulfate release from green rust during the sorption experiments was between 416 and 625 µmols SO 4 /g green rust using initial phosphate concentrations between 10 and 1500 µm and was not directly correlated to PO 4 (Figure 2.6). The NaCl used to adjust the ionic strength of the phosphate solutions, however, had an appreciable influence on sulfate release (Figure 2.5). Green rust allowed to react with anoxic distilled water for two days released an average of µmols mg SO 4 /g and that reacted with 0.1 M NaCl released µmols SO 4 /g. When a 100 µm NaH 2 PO 4 solution (without the ionic strength adjusted with NaCl) was used about µmols SO 4 /g green rust was released. Additionally, green rust reacted with 0.1 M NaCl and 1500 µm NaH 2 PO 4 released the same amount of sulfate per gram as green rust treated with only 0.1 M NaCl. The 55

24 presence of chloride is not only the predominant factor in how much sulfate is released, but also is a factor in how much phosphate will be sorbed. Green rust was reacted with 100 µm NaH 2 PO 4 solutions with and without the ionic strength adjusted; the samples without NaCl sorbed more phosphate than those reacted with a solution containing 0.1 M NaCl (66.3 vs µmols PO 4 /g green rust, respectively). These observations indicate a preference by interlayer sulfate to exchange with chloride over phosphate. This preference may only be due to the concentration gradient betweens Cl - and PO 3-4. Lewis (1997) found that chloride green rust would convert to sulfate green rust when reacted with solutions containing less than 100 mm SO 4 ; however, the study did not examine whether the reverse reaction would occur. Generally, a decrease in adsorption is seen with increasing ph; however, this decrease is minimal for the ferric oxides within the ph range found in most soils (4 7) (Brady and Weil, 1999). The trends displayed by the goethite- and ferrihydrite-coated sand are in agreement with the findings of Geelhoed (1997), Strauss et al. (1997), Goldberg (1983), Willett and Cunningham (1983), Ryden et al. (1977a), and McLaughlin et al. (1977). The maximum sorption on goethite and ferrihydrite may not have been observed and may actually occur at ph near 2.5 (Strauss et al., 1997). Both magnetite samples and the natural siderite had well defined adsorption envelopes, with adsorption rapidly diminishing between ph 8 and 9 (Figure 2.4). The ferrihydrite and ferrihydritecoated sand behave differently with increased ph. Adsorption onto the coated sand decreases much more rapidly with ph than does sorption onto the ferrihydrite slurry. The calculated binding coefficients were for the ferrihydrite slurry and for the coated sand. Since the ferrihydrite slurry has a greater affinity for phosphate, one would expect the changes in phosphate adsorption with ph to be less dramatic than those of the ferrihydrite-coated sand. The trends in decreasing adsorption with increasing ph will have implications on the fate and transport of phosphate in natural systems. When phosphate is adsorbed to a ferric hydroxide, two hydroxyl groups are released to solution, resulting in an initial removal of phosphate from a system along with an increase in ph. This ph increase will slow the reaction kinetics of sorption (Shang et al., 1992) and even completely inhibit adsorption - as observed in the case of both magnetite samples and the geologic siderite. 56

25 Inhibition of phosphate sorption may occur at ph higher than what was examined in this study for other iron oxides. Hingston et al. (1972) showed that phosphate sorption onto goethite decreased to zero at ph 13. In terms of practical application, if a permeable reactive barrier utilizing iron oxides were to be used for phosphate removal, one would need to consider ways to control the ph. In a system that is not well buffered, the initial phosphate adsorption will lead to increased ph and subsequently inhibited phosphate removal. The retention would be slowed or completely halted until the ph decreased again; in either case, the result would be ineffective phosphate removal. However, in most soil systems the buffering capacity is likely to be great enough to mitigate large ph changes upon P sorption. The sorption characteristics of the siderite and magnetite samples varied considerably between the geologic and synthetic specimens. In the case of the magnetite, both exhibited a well-defined adsorption envelope, but the synthetic magnetite reached an apparent adsorption maximum and was fit with a Langmuir isotherm; the geologic sample did not exhibit an adsorption maximum and was fit to the Freundlich equation. Adsorption onto synthetic siderite did not appear to change dramatically with ph; however, sorption by the natural sample decreased to insignificant amounts at ph 9. Similarly to the magnetite, no adsorption maximum on geologic siderite was observed and was fit with the Freundlich equation, whereas phosphate adsorption onto synthetic siderite conformed to the Langmuir equation. Additionally, dissolved iron was detected in phosphate solutions following a two-day reaction period with both natural samples. These variations may be the result of a variety of factors. One consideration is mineral purity; the synthetic specimens are likely to have a much more consistent composition. Geologic magnetite is commonly found to have small amounts of Ca 2+, Mn 2+, and Mg 2+ replacing Fe 2+, and Al 3+ can replace Fe 3+. Other, less common, replacement elements include Cr 3+ and V 3+ (for Fe 3+ ) and Ni 2+, Co 2+, and Zn 2+ (for Fe 2+ ). Manganese and magnesium are both common substitutions for Fe 2+ in siderite since there is complete solid solution between siderite and rhodochrosite and siderite and magnesite (Deer et al., 1992). Ainsworth et al. (1985) found that Al-substitution in goethite had an appreciable influence on the extent and rate of isotopic phosphate exchange. The siderite and magnetite samples were treated with dilute acid prior to the experiments; however, 57

26 residual oxidative coatings and inclusions may have influenced the sorption behavior. Magnetite has commonly been observed to convert to maghemite and hematite (Cornell and Schwertmann, 1994), and magnetite crystals are often observed to contain detectable amounts of maghemite (Deer et al., 1992). Siderite usually will alter to goethite and sometimes hematite (Deer et al., 1992). In general, phosphate adsorption changed very little over the range of ionic strengths examined. Sorption by natural magnetite and ferrihydrite (slurry) decreased slightly with ionic strength and the remaining minerals, with the exception of green rust, increased in sorption capacity with ionic strength. Green rust, discussed previously, exhibited a substantial decrease in phosphate adsorption with increased chloride concentration. Ionic strength alone is likely to have less influence on phosphate retention than the ph and substrate; however, competition of other ions in solution for binding sites is a factor that must be considered. Ryden et al. (1977b) found that increased ionic strength resulted in an increase in phosphate sorption in soils and that the increase was more pronounced with Ca present than with Na in solution. Additionally, ph may influence how dramatically fluctuations in ionic strength affect phosphate sorption. Arai and Sparks (2001) found that from ph 4 to 7.5, phosphate adsorption on ferrihydrite was unaffected by changes in ionic strength; however, at ph greater than 7.5, adsorption increased with ionic strength. Patrick and Khalid (1974) postulated that the reductive dissolution of ferric iron minerals led to ferrous compounds having increased activities and surface areas but that the ferric oxides had a higher affinity for phosphate. These postulates are supported by the findings of our work. On a surface area basis, all the ferrous iron minerals had much higher capacities to sorb phosphate (from to µmols PO 4 /g mineral) than the ferric minerals (1.68 to 6.32 µmols PO 4 /g mineral). This would result in the increased phosphate sorption capacity of reduced sediments and soils in many areas provided high surface area phases developed; however, the calculated binding and distribution coefficients indicate that the ferric minerals have a higher affinity for phosphate than do the ferrous minerals. The findings of this research also help to explain other previous observations and have implications in how excess phosphate could be removed from a system. Patrick and Khalid (1974) also noted that even though anaerobic soils were 58

27 capable of sorbing more phosphate they also released more into the soil solution, an observation believed to result from lower affinity for phosphate by the ferrous minerals. The distribution and binding coefficients calculated during this work support that theory (Table 3). The adsorption characteristics exhibited by the ferric and ferrous iron minerals also help to explain why seasonally flooded soils have been observed to have higher levels of non-labile (non-exchangeable/iron bound) phosphate (Reynolds et al. 1999). When flooded soils are subsequently aerated the relatively fine-grained ferrous minerals would re-oxidize resulting in ferric phases with higher surface area and thus more sorption sites than the ferric minerals present prior to flooding. The higher affinities of the ferric iron minerals result in more strongly held (less labile) phosphate in the soil. Changes in ph have been observed to affect phosphate retention by numerous ferric oxides (Geelhoed, 1997; Strauss et al., 1997, Shang et al., 1992; Willett and Cunningham, 1983; Ryden et al., 1977b; McLaughlin et al., 1977). The adsorption envelopes observed in this study indicate ph could play a much larger role in influencing phosphate retention in sub-oxic to anoxic soils than in well-aerated systems. CONCLUSIONS The results of this study indicate ferrous minerals have phosphate adsorption capacities that are much greater than those of the ferric phases examined. However, the ferrous minerals all had much lower binding coefficients and were more sensitive to changes in ph. The high capacity for phosphate of the ferrous minerals, combined with the relatively low affinity and sensitivity to ph suggests that the ferrous minerals investigated may provide appreciable, although transitory phosphorus sinks in some reductomorphic environments. 59

28 CHAPTER 2 APPENDIX: ISOTHERM DATA COLLECTED DURING SORPTION EXPERIMENTS 60

29 Table 2A.1. Summary of data used to establish adsorption isotherm of PO 4 by naturally occurring (geologic origin) magnetite. Sample PO 4 solution (L) g magnetite M P (end) M P (start) Change in M P M P/g magnetite mg PO 4 /g magnetite Avg mg/g SD mg/g SA m 2 /g M/m 2 magnetite Avg M/ m 2 SD M/m 2 10A B C A B C A B C A B C A B C A B C A B C A B C ph 4 A n/a ph 4 B ph 4 C ph 5.5 A n/a ph 5.5 B ph 5.5 C ph 8 A ph 8 B ph 8 C ph 9 A ph 9 B ph 9 C I/S A I/S B I/S C

30 Table 2A.2. Summary of data used to establish adsorption isotherm of PO 4 by sulfate green rust. Sample PO 4 solution (L) g green rust M P (end) M P (start) Change in M P M P/g green rust mg PO 4 /g green rust Avg mg/g SD mg/g SA m 2 /g M/m 2 green rust Avg M/ m 2 SD M/m 2 10A B A B A B A B A B A B A B A B ph 4 A ph 4 B ph 5.5 A ph 5.5 B ph 8 A ph 8 B ph 9 A ph 9 B I/S A I/S B