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Soil Science Society of America Journal 64:885-892 (2000)
© 2000 Soil Science Society of America

DIVISION S-2-SOIL CHEMISTRY

Perturbation of Taranakite Formation by Ferrous and Ferric Iron under Acidic Conditions

J.M. Zhoua,b, C. Liua and P.M. Huanga

a Dep. of Soil Science, Univ. of Saskatchewan, 51 Campus Drive, Saskatoon, SK Canada S7N 5A8
b The Institute of Soil Science, Academia Sinica, P.O. Box 821, 71 East Beijing Road, Nanjing 210008, China

huangp{at}sask.usask.ca


   ABSTRACT
TOP
 ABSTRACT
INTRODUCTION
 Materials and methods
 Results and discussion
 Conclusions
 REFERENCES
 
Taranakite is an important reaction product of monoammonium phosphate fertilizer with soils. Its formation affects the transformation of nutrients in soils. The effect of different molar ratios of Fe(II)/Al and Fe(III)/Al on the formation of taranakite at pH 4.0 was investigated in this study. The results show that Fe(II) ion significantly perturbed the formation of taranakite at the Fe/Al molar ratio of 1.2 . When the Fe/Al molar ratio was increased to 2.5, the formation of taranakite was completely inhibited by Fe(II), whereas, under the same condition, some crystalline taranakite was still observed in the Fe(III) system. Although Fe(III) had less effect on the crystallization of taranakite than Fe(II) at lower Fe/Al molar ratios, it also completely inhibited the formation of taranakite at the molar ratio of Fe(III)/Al >= 5. The solid products formed in the Fe(III) or Fe(II) system contained a substantial amount of Fe(III) and a much higher proportion of phosphate than that was required for the formation of NH4–taranakite. As indicated by the solution phase analysis at the end of the experiment, more Fe ions were present in the solution in the Fe(II) system, compared with the Fe(III) system, to perturb the nucleation and crystallization of taranakite. Since iron is a very common element in soil, taranakite formation may be perturbed in soils with high Fe content, especially under reduced and acidic conditions.

Abbreviations: XRD, x-ray diffraction • IR, infrared absorption


   INTRODUCTION
TOP
 ABSTRACT
 INTRODUCTION
Materials and methods
 Results and discussion
 Conclusions
 REFERENCES
 
TARANAKITES, including ammonium taranakite [(NH4)3Al5H6(PO4)8·18H2O] and potassium taranakite [K3Al5H6(PO4)8·18H2O], have the same structure (Frazier and Taylor, 1965), although their composition is different. Natural K–taranakite was first found from trachytic rocks of the Sugarloaves, Taranaki, New Zealand in 1865 (Bannister and Hutchinson, 1947). This mineral was later found in some caves of France, Italy, Algeria, Australia, Reunion (an island of the Indian Ocean), USA, and Japan (Bannister and Hutchinson, 1947; Murray and Dietrick, 1956; Sakae and Sudo, 1975; Fiore and Laviano, 1991).

Application of fertilizers, which contain phosphate, ammonium, and potassium, may result in the formation of taranakites as reaction products in the immediate vicinity of fertilizer zones. Lindsay et al. (1962) used saturated solutions of phosphate fertilizers to react with soils for various periods to simulate the reaction zone of fertilizer bands. After the reaction, the filtrates of the suspensions were stored in the laboratory for periods from 1 d to several months to await precipitation of the solid phase. The formation of taranakites was observed in the precipitates of the monoammonium phosphate and monopotassium phosphate systems. By using the same method, taranakites were also found by other researchers (Sarkar et al., 1977; Prabhudesai and Kadrekar, 1984). Recently, ammonium taranakite was directly observed in 1 M NH4H2PO4–treated soils (Zhou and Huang, 1995).

Taranakites contain P, N, and K in their structures. An evaluation of taranakites as a source of phosphate for plants was conducted by Taylor et al. (1960). The formation of taranakites in soil can result in the transformation of these nutrients from the readily available form to the slowly available form and, thus, affect the dynamics and bioavailability of the nutrients in soil to plants, especially in the immediate vicinity of phosphate fertilizer bands. This transformation may reduce the rate of P, N, and K supply at the early stage of plant growth but may be beneficial to the long-term supply of the nutrients to plants.

The application of phosphate fertilizers may dissolve soil minerals, resulting in the release of Al ions and the subsequent formation of taranakites. However, other cations and anions which coexist with Al in soil minerals may also be dissolved. This could perturb the formation of taranakites and, thus, influence the dynamics and bioavailability of P, N, and K in soils. Iron is one of the most abundant elements in soil environments. It may exist in ferrous or ferric forms depending on the composition of the soil minerals and the redox condition of the soil. According to the previous study (Zhou and Huang, 1999), the pH range for the formation of NH4–taranakite is from 2.75 to 5.75. The objective of this research was to investigate the effect of ferrous and ferric iron on the formation of taranakite in acidic condition similar to phosphate fertilizer zones.


   Materials and methods
TOP
 ABSTRACT
 INTRODUCTION
 Materials and methods
Results and discussion
 Conclusions
 REFERENCES
 
Synthesis of Taranakite
In this study, only NH4–taranakite was studied because ammonium phosphate fertilizer is more commonly used than potassium phosphate. Ammonium taranakite can be synthesized by using NH4H2PO4 and AlCl3. The NH4H2PO4 and AlCl3 used in the study were pure chemicals (Analar grade, BDH Inc., Toronto, ON, Canada). In order to simulate the condition of a phosphate fertilizer band in soil, a high concentration of NH4H2PO4 (1 M) was used in the synthesis. The concentrations of AlCl3 used were 1 x 10-2 and 5 x 10-2 M. After mixing 50 mL of a 2 M NH4H2PO4 solution and 10 mL of a 0.1 M AlCl3 solution in a 100 mL Erlenmeyer flask, the mixed solution was diluted to close to 100 mL by using deionized distilled water. The pH of the solution was adjusted to 4.0 by using ammonia water with continual stirring. A white precipitate appeared during the pH adjustment. The suspension was then made to exactly 100 mL. The final concentrations of NH4H2PO4 and AlCl3 were 1 and 1 x 10-2 M, respectively. For the 5 x 10-2 M AlCl3 system, the same procedure was used except that, instead of 10 mL of 0.1 M AlCl3 solution, 5 mL of 1 M AlCl3 solution was added to the system. The container that stored the suspension was sealed by Parafilm and stored in the laboratory at 23°C. After 1 d to 4 wk, the suspension was filtered through a Millipore membrane with a pore size of 0.01 µm. Two milliliters of concentrated HCl were immediately added to the filtrate to prevent Fe(II) from oxidation (Tamura et al., 1976). The solid phase on the membrane was washed with deionized distilled water until no Cl- in the filtrate was detected with a 1% AgNO3 solution. The sample was then air dried for identification of the reaction products. As a reference for comparison, NH4–taranakite was also synthesized in the present study according to the method described by Taylor and Gurney (1961).

Iron Perturbation of Taranakite Formation
In the synthesis process of taranakite, FeCl2 or FeCl3 was mixed with AlCl3 before the AlCl3 solution was added to the NH4H2PO4 solution. Two concentrations of Al (1 x 10-2 and 5 x 10-2 M) were used in this study. For each concentration of Al, the molar ratios of Fe/Al were adjusted to 0, 1.2, 2.5, and 5.0, respectively. The pH was adjusted to 4.0, which was the same as the pure taranakite systems. The FeCl2 solution was freshly prepared immediately before the experiment. To minimize Fe(II) oxidation during the experiment, a strong reducing agent (0.15% N2H4 with a purity of 98% ) was added to the Fe(II) systems to compare with the Fe(III) system and the Fe(II) system without N2H4. In the N2H4 treatment, the suspension was flushed with N2, the container was sealed with Parafilm, and the suspension was aged for 3 wk. The other procedures involved in the study of perturbation of iron on the formation of taranakite were the same as for the synthesis of taranakite.

Identification of the Solid Phases
Fifteen milligrams of each washed and air-dried sample were lightly ground and mounted on a glass slide by adding a few drops of acetone without any further treatment prior to the x-ray diffraction (XRD) analysis. The analysis was carried out using a Philips (Model PW 1031) x-ray diffractometer (Eindhoven, the Netherlands) using Mn-filtered Fe-K{alpha} radiation at 35 kV and 16 mA. For infrared absorption (IR) analysis, 2 mg of each sample were mixed with 250 mg KBr and then pressed into discs. Using a KBr disc as a reference, the discs with samples were examined with a Perkin-Elmer (Model 983) infrared absorption (IR) spectrophotometer (Buckinghamshire, UK).

Determination of pH, Eh, and Element Concentrations
The initial and final pH and Eh values of each treatment were determined. The pH of suspensions was measured by using a glass–calomel combination electrode. The Eh was determined by using a Pt electrode on a Metrohm titroprocessor (683 model, Metrohm Ltd., Herisau, Switzerland). The concentrations of P, NH+4–N, Al, Fe(III), and Fe(II) in the acidified filtrate solutions were determined. The Al and total Fe concentration in the solutions were, respectively, measured using atomic absorption spectrometry at 309.3 and 248.3 nm on a Perkin Elmer atomic absorption spectrophotometer (Model 3100, Perkin Elmer Corp., Norwalk, CT). The ferrous Fe concentration was measured by the colorimetric method with 2,4,6-tri(2'-pyridyl)-1,3,5-triazine (Krishnamurti and Huang, 1990). The ferric Fe concentration was calculated based on the difference between the total Fe concentration and ferrous Fe concentration. The P concentration was determined by the vanadomolybdophosphoric acid colorimetric method (Kuo, 1996). The NH+4–N concentration was measured by the emerald green colorimetric method (Mulvaney, 1996). The contents of P, NH+4–N, Al, Fe(III), and Fe(II) in the solid phases were determined by dissolving 0.02 g of products in 2 mL of concentrated HCl, which was then diluted to 100 mL. The concentration of P, NH+4–N, Al, Fe(III), and Fe(II) were determined by the same methods as described above.

The experiment in the present study was carried out in triplicate.

Calculation of Species Using Geochemical Modeling Program
The chemical species in each system of taranakite formation, except for the Fe(III) system, were calculated by using geochemical modeling program FITEQL (version 3.1) (Herbelin and Westall, 1994). The ionic strength of the system in the present study ranged from 1.05 to 1.48. The Davies equation is implemented to compute activity coefficients in the FITEQL program. However, the Davies equation is not adequate for high ionic strength solutions. Therefore, when the FITEQL program was used to model the chemical species in this study, the activity coefficient calculation was thus set to NO. The concentrations of Al, Fe(II), phosphate, and NH+4 were converted to activities based on the activity coefficient calculated from Pitzer equation (Pitzer, 1991) and then the activities were input into the FITEQL program. Since the parameters ß and C{Phi} (respectively for the binary and tertiary solute-solute interactions) for FeCl3 are not available, the chemical species in the Fe(III) system were not calculated. The pertinent equilibrium constants for reactions which were considered in the modeling are listed in Table 1 . Since Fe(II) ion was continuously oxidized to Fe(III) during the reaction and aging processes, the Fe(II) concentration changed with time. The chemical species of taranakite formation system in the presence of Fe(II) were modeled by two approaches, that is, including and excluding the redox reaction: . The two approaches respectively simulated the reactions occurring under the final and initial conditions of the system.


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  		Table 1 Equilibrium reactions considered in the chemical modeling.{dagger}

 

   Results and discussion
TOP
 ABSTRACT
 INTRODUCTION
 Materials and methods
 Results and discussion
Conclusions
 REFERENCES
 
Effect of Iron(II) and Iron(III) on the Formation of Taranakite at Different Iron/Aluminum Molar Ratios
The NH4–taranakite synthesized in this experiment was compared with the NH4–taranakite synthesized by the method given by Taylor and Gurney (1961). The samples synthesized by both methods had the same XRD patterns and IR spectra, indicating the reliability of the synthesis method used in this study.

Based on the solubility of aluminum hydroxides which are common in soils, such as amorphous Al(OH)3 and gibbsite, the activity of Al3+ ions at pH 4.0 is in the range of about 5 x 10-3 to 2 x 10-4 (activity is unitless) (Lindsay, 1979). The initial Al concentration in the present study was in the range of 10-2 to 5 x 10-2 M, corresponding to the Al3+ activity of 7.23 x 10-5 to 1.92 x 10-4 in the presence of 1 M NH4H2PO4. The activity coefficients used to convert the activities to the concentrations were computed based on the Pitzer equation (Pitzer, 1991). The initial Fe(II) concentration in this study ranged from 1.2 x 10-2 M to 6 x 10-2 M that is equivalent to Fe2+ activity of 9.72 x 10-4 to 6.58 x 10-3 in the presence of 1 M NH4H2PO4. According to the diagram of the solubility relationships of iron with pe + pH of typical soils (Schwab and Lindsay, 1983), the Fe2+ activity controlled by goethite at the pe + pH of 6, which was the initial pe + pH of the Fe(II) system in this study, is about 10-1. Thus, the initial Al3+ and Fe2+ activities used in the present study were in the range of Al3+ and Fe2+ activities in common soil solutions. Since the parameters ß and C{Phi} for FeCl3 are not available, the Fe3+ activity in the Fe(III) system cannot be calculated according to the Pitzer equation.

The x-ray diffractograms of the precipitates formed under the conditions of different Fe(II)/Al ratios are presented in Fig. 1 . Compared with the pure taranakite (Fig. 1a), Fe(II) did perturb the formation of taranakite and the perturbation became more pronounced as the molar ratio of Fe(II)/Al increased (Fig. 1b, 1c, and 1d). At the end of the aging period of 4 wk, only noncrystalline products were formed when the Fe(II)/Al molar ratio was 2.5 or higher.



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  		Fig. 1 X-ray diffractograms of the precipitates formed at 1 M NH4H2PO4, 10-2 M Al, pH 4.0, and various Fe(II)/Al molar ratios. The aging period was 4 wk at 23°C

 
The effect of Fe(III) on the formation of taranakite is shown in Fig. 2 . At a Fe(III)/Al molar ratio of 1.2, the XRD pattern (Fig. 2b) was very similar to the control (Fig. 2a). Even at the Fe(III)/Al molar ratio of 2.5, the crystalline taranakite was still observed in the XRD pattern (Fig. 2c) although the perturbation by the Fe(III) was clear. Ferric iron completely inhibited the formation of taranakite at the Fe(III)/Al ratio of 5.0 (Fig. 2d). Therefore, both Fe(II) and Fe(III) can affect the formation of taranakite, but the influence of Fe(II) was substantially stronger than Fe(III).



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  		Fig. 2 X-ray diffractogram of the precipitates formed at 1 M NH4H2PO4, 10-2 M Al, pH 4.0, and various Fe(III)/Al molar ratios. The aging period was 4 wk at 23°C

 
Figure 3 shows the x-ray diffractograms of the precipitates formed in Fe(II) and Fe(III) systems at the Fe/Al ratio of 1.2 after aging for 3 wk. In the presence of Fe(III), the XRD pattern (Fig. 3b) clearly shows the formation of crystalline taranakite. However, virtually only noncrystalline products were formed in the presence of Fe(II) (Fig. 3c and 3d). The IR spectra of the same precipitates are presented in Fig. 4 . The IR spectrum of pure NH4–taranakite was shown by Frazier and Taylor (1965). The phosphate absorption bands were from 887 to 1193 cm-1, whereas ammonium absorption bands were at 1430, 1461, and 3240 cm-1. The absorption bands of 1653 and 3110 cm-1 were from H2O. The 2422 cm-1 band may represent the H–O–P bond, whereas the bands from 328 to 606 probably were related to the Al–O–P bond (Fig. 4a).



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  		Fig. 3 X-ray diffractograms of the precipitates formed at 1 M NH4H2PO4, 10-2 M Al, pH 4.0, and various Fe/Al molar ratios under different redox conditions: (a) , (b) Fe(III), , (c) Fe(II), , and (d) Fe(II) +N2H4, . The aging period was 3 wk at 23°C

 


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  		Fig. 4 Infrared spectra of the precipitates formed at 1 M NH4H2PO4, 10-2 M Al, and pH 4.0 in different iron oxidation state system. The aging period was 3 wk at 23°C

 
In the presence of Fe(II), some absorption bands in taranakite disappeared or shifted (Fig. 4c). The absorption bands from 328 to 606 cm-1 were changed to two bands at 460 and 538 cm-1. The phosphate absorption bands from 887 to 1193 cm-1 were changed to a single band of 1058 cm-1. The ammonium absorption bands at 1430 and 1461 cm-1 were transformed to 1399 cm-1, whereas the 3240 cm-1 band from ammonium was shifted to 3185 cm-1. The appearance of the broad absorption bands indicated the formation of noncrystalline products. Compared with Fe(II), the degree of perturbation for the formation of taranakite by Fe(III) was much less (Fig. 4b and 4c). The IR data were in accord with the results of XRD (Fig. 3).

Since a high concentration of NH4H2PO4 (1 M) and a relatively low concentration of AlCl3 (10-2 M)were used in the formation of taranakite, the concentrations of P and N remaining in the solution after a 4-wk aging period were only slightly decreased (Table 2) . However, the concentration of Al remaining in the solution was decreased by two to three orders of magnitude compared with the initial Al concentration (Table 2). In the presence of Fe(II) or Fe(III) ions, more Al ions remained in the solution at the end of a 4-wk aging period compared with the system in the absence of Fe, except for the Fe(III)/Al molar ratio of 1.2 at the initial Al concentration of 10-2 M (Table 2). In this case, Al ions remained in the solution virtually at the same concentration as the Fe/Al molar ratio of 0. The x-ray patterns (Fig. 2) show that well-crystalline taranakite was formed at the Fe(III)/Al molar ratios of 0 and 1.2. In the Fe(II) systems, no Fe(II) ions were detected in the solution at the end of a 4-wk aging period (Table 2), indicating that Fe(II) ions were eventually oxidized to Fe(III) ions using O2 as an electron acceptor. In the Fe(II) system, less than 35% of the initial Fe ions was present in the solutions as Fe(III) ions at the end of a 4-wk aging period. At the same Fe/Al molar ratios, much fewer Fe ions remained in the solution in the Fe(III) system compared with the Fe(II) system (Table 2), indicating that more Fe ions were precipitated in the Fe(III) system.


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  		Table 2 The concentration of Al, P, NH+4-N, Fe(III), and Fe(II) remaining in solutions in the taranakite systems after aging for 4 wk

 
Elemental composition of the solid products is given in Table 3 . Based on the formula [(NH4)3Al5H6 (PO4)8·18H2O], NH4–taranakite contains 105.4 g Al kg-1, 193.7 g P kg-1, and 32.8 g N kg-1. Compared with the formula of NH4–taranakite, the precipitates formed in the absence of Fe ions had the same P content and about 13% lower Al and 4% lower N contents (Table 3). The lower Al and N contents of the taranakite formed probably resulted from crystal defects and/or the presence of some noncrystalline precipitates. In the 10-2 M Al system, the Al and N contents of the products formed in the presence of Fe(II) or Fe(III) drastically decreased compared with those of the products formed in the absence of Fe (Table 3). The higher the Fe(II)/Al or Fe(III)/Al molar ratio, the lower were the Al and N contents.


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  		Table 3 Elemental composition of the solid products in the taranakite systems after aging for 4 wk

 
The P content of the products formed in the presence of Fe(II) or Fe(III) also decreased compared with taranakite formed in the absence of Fe (Table 3). There was no detectable Fe(II) in the solid phase products in the Fe(II) system but substantial amounts of Fe(III) were present in the solid phase products (Table 3). Fe(II) was evidently oxidized to Fe(III) in the precipitates during the reaction processes and aging period. At the same Fe/Al molar ratio, the Al, N, and P contents of the products formed in the Fe(III) systems were substantially lower than those formed in the Fe(II) systems (Table 3).

The P/Al/N ratio of the taranakite formed in the absence of Fe ions was 6.14:2.91:1. The similar P/Al/N ratio should have been observed if taranakite were the reaction product in the solid phase in the presence of Fe(III) or Fe(II). At the Fe(III)/Al molar ratio of 1.2, the product had the Al/N ratio of 2.90:1. The XRD pattern also shows that well-crystalline taranakite was formed at this Fe(III)/Al molar ratio (Fig. 2b). However, this product had a much higher P/N ratio than the taranakite formed in the absence of Fe (Table 3). The excess P presumably coprecipitated with Fe(III), since a large amount of Fe (125 g Fe kg-1) were detected in the product. The XRD pattern of the products only shows the d-values of taranakite. Therefore, the iron phosphate formed was x-ray noncrystalline.

When the Fe(III)/Al molar ratio was increased to 2.5 and 5.0, the Al/N ratio of the products formed drastically decreased (Table 3), indicating that the formation of taranakite was perturbed. This is corroborated by the XRD data (Fig. 2c and 2d). Compared with the solid phases formed at the Fe(III)/Al molar ratio of 1.2, the higher P/N ratio and larger amounts of Fe of the products formed at the Fe(III)/Al molar ratios of 2.5 and 5.0, along with the XRD data indicate that more noncrystalline iron phosphates were formed. These iron phosphate phases may interfere with the taranakite formation by affecting the crystal growth during the crystallization.

At the same Fe/Al molar ratios, the Al/N ratio of the products formed in the Fe(II) system was lower than that of the products formed in the Fe(III) system (Table 3), indicating that the formation of taranakite was more perturbed in the Fe(II) system than in the Fe(III) system. This is consistent with the XRD data (Fig. 1b, 1c, 1d, 2b, 2c, and 2d). Compared with the Fe(III) system, the lower P/N ratio of the products formed in the Fe(II) system at the same Fe/Al molar ratio indicates that less iron phosphate phases would be formed, although the P/N ratio was higher than that of the taranakite formed in the absence of Fe (Table 3). This is in accord with the higher concentration of Al and Fe remaining in the solution in the Fe(II) system compared with the Fe(III) system (Table 2). However, the solid products formed in the Fe(II) system had a very similar Fe content compared with those formed in the Fe(III) system (Table 3). Apparently, Fe more preferentially coprecipitated with Al to perturb the formation of taranakite in the Fe(II) system than in the Fe(III) system. The substitution of some Al by Fe(III) was found in the natural K–taranakite sample at Taranaki, New Zealand (Bannister and Hutchinson, 1947).

The results of chemical modeling (Table 4) along with the experiment data on the concentration of Al, P, N, Fe(II), and Fe(III) remaining in the solution in the aging period of 1 d to 4 wk (data not shown) indicate that the experimental system had not completely reached equilibrium. In the Fe(II) system, the aqueous species were modeled by two approaches. In the first approach, the oxidation of Fe(II) to Fe(III) was not considered to simulate the initial reaction system in which most of Fe(II) was not oxidized. In this case, the activity of FeH2PO+4 complex was higher than that of free Fe2+ and its hydrolyzed species (Table 4). Since Fe(II) was continuously oxidized, the Fe(II) activity would be steadily decreased. Therefore, in the second approach, the oxidation of Fe(II) to Fe(III) was considered so as to simulate the final reaction system based on the final pe + pH value. In this case, the Fe ions should remain in the solution as both Fe(II) and Fe(III) (Table 4). However, the Fe(II) was not detected in the experiment (Table 2). This is because the presence of O2 continually promoted the oxidation of Fe(II) to Fe(III) as discussed before. The activity of complexes of Fe(III) and Fe(II) with phosphate was much higher compared with the activity of the free Fe ions and their hydrolysis species (Table 4). This indicates that iron perturbed taranakite formation dominantly through the formation of the phosphate complexes rather than the hydrolysis species.


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  		Table 4 The major species of aqueous phases in the taranakite systems based on chemical modeling.{dagger}

 
The experimental data show that the solution Fe concentration in the Fe(II) system was one to two orders of magnitude higher than that in the Fe(III) system at the end of a 4-w aging period (Table 2). Therefore, the collision of phosphate, ammonium, Al, and H ions to form the nuclei of taranakite by diffusion was much more affected by the solution Fe in the Fe(II) system than in the Fe(III) system. This explains the observation that the effect of Fe(II) on the formation of taranakite was stronger than Fe(III). Furthermore, the solution Fe activity (all solution species of Fe) calculated by the first approach which excludes the oxidation of Fe2+ to Fe3+ was much higher than that calculated by the second approach which includes the oxidation of Fe2+ to Fe3+ (Table 4). This indicates that the perturbation of taranakite formation by Fe(II) was more pronounced at the initial stage of the reaction before the Fe2+ was oxidized to Fe3+.

Effect of Eh and Aging Period on the Perturbation of the Formation of Taranakite by Iron
The initial and final Eh values and the final pH values of the taranakite formation systems are given in Table 5 . The final pe + pH values in the present study ranged from 9.45 to 10.20 which are very common in soils. The final Eh of the Fe(II) system was virtually the same as that of the Fe(III) system, although their initial Eh values were very different. During the experiment, oxidation of Fe(II) to Fe(III) (Tables 2 and 5) weakened the effect of Fe on the formation of taranakite. In order to prevent Fe(II) from being oxidized, a strong reducing agent (N2H4) was added to the Fe(II) system at the beginning of the experiment. The addition of N2H4 resulted in the decrease in both initial and final Eh values in the Fe(II) system. The x-ray diffractograms of the precipitates formed under these different conditions are shown in Fig. 3. In the Fe(II) system without N2H4, two weak peaks (16.33 and 8.06 ) were found in the XRD pattern (Fig. 3c), indicating the formation of a very small amount of taranakite. In the Fe(II) system with N2H4, the formation of taranakite was completely inhibited (Fig. 3d). The difference between the Fe(II) and the Fe(II) + N2H4 systems was due to the oxidation of Fe(II) in the system in the absence of the reducing agent. In the Fe(II) + N2H4 system, more Fe(II) would be present in the solution to perturb taranakite formation.


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  		Table 5 Initial and final Eh values and final pH values in the taranakite systems after aging for 4 wk

 
Although only weak peaks of crystalline taranakite were observed in the Fe(II) system after aging for 3 wk (Fig. 3c), the peak intensity of 16.33 and 8.06 in the precipitate increased as the aging period increased to 4 wk (Fig. 1b). The formation of more crystalline taranakite during the longer aging period was attributed to the recrystallization process of the precipitates.

Effect of Aluminum Concentration on the Perturbation of the Formation of Taranakite by Iron(II) and Iron(III)
The concentration of Al in soil solution varies with soil environments. This variation could lead to the change in sensitivity of the perturbation of the taranakite formation by Fe. The x-ray diffractograms of the precipitates formed in different concentrations of Al at the Fe/Al molar ratio of 1.2 are presented in Fig. 5 .



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  		Fig. 5 X-ray diffractograms of the precipitates formed at 1 M NH4H2PO4, pH 4.0, an Fe/Al molar ratio of 1.2, and different Al concentrations: (a) Fe(II), , (b) Fe(II), , (c) Fe(III), , and (d) Fe(III), . The aging period was 3 wk at 23°C

 
At an Al concentration of 5 x 10-2 M, well crystalline taranakite was formed and the Fe perturbation of taranakite formation was not as significant in either the Fe(II) or Fe(III) system (Fig. 5a and 5c), compared with the lower Al concentration systems (Fig. 5b and 5d). At the initial Al concentration of 5 x 10-2 M, more Al and Fe ions and less P and NH+4 ions remained in the solution compared with the reaction system at the initial Al concentration of 10-2 M (Table 2). The Al/N ratio of the products formed in the presence of either Fe(III) or Fe(II) at 5 x 10-2 M Al was very similar to that of the product formed in the absence of Fe (Table 3), indicating the formation of taranakite. However, the P/N ratio of the products formed in the presence of either Fe(III) or Fe(II) was much higher than that of the product formed in the absence of Fe (Table 3). This along with the Fe contents of the solid products (Table 3) and the XRD data (Fig. 5) indicates that an x-ray noncrystalline iron phosphate was formed with taranakite.

In the systems studied, the concentrations of ammonium and phosphate were very high (1 M), therefore the Al concentration was one of the critical factors in determining the rate of taranakite formation. When the Al concentration was increased, the rate of nuclei formation of taranakite increased rapidly, because the probability of a simultaneous collision of a large number of particles depends on the concentration of reaction species (Khamskii, 1969). Therefore, the higher Al concentration led to the easier nucleation and crystallization of taranakite. If the Al concentration was lower in the system, the nucleation and crystallization of taranakite were slow; the Fe in the system, thus, had more chance to perturb the formation of taranakite through the complexation with phosphate. Therefore, the lower the concentration of Al in the system, the more the perturbation in taranakite formation by Fe was observed.


   Conclusions
TOP
 ABSTRACT
 INTRODUCTION
 Materials and methods
 Results and discussion
 Conclusions
REFERENCES
 
Both Fe(II) and Fe(III) perturbed the formation of taranakite at pH 4.0 and at the Al concentration of 10-2 to 5 x 10-2 M. The extent of the perturbation increased as the Al concentration decreased and the molar ratio of Fe/Al increased. Ferrous iron was more effective in perturbing the formation of taranakite than Fe(III). The solution analyses show that more Fe ions were present in the solution to affect the nucleation and crystallization of taranakite in the Fe(II) system compared with the Fe(III) system. Therefore, the role of Fe in perturbing the formation of crystalline taranakite in the immediate vicinity of fertilizer zones in soils, especially under reduced conditions, and the impact on the dynamics and bioavailability of P and N merit close attention.


   ACKNOWLEDGMENTS
 
This study was supported by a grant of the Potash and Phosphate Institute of Canada and Grant GP 2383-Huang of Natural Sciences and Engineering Research Council of Canada.

Received for publication October 2, 1998.


   REFERENCES
TOP
 ABSTRACT
 INTRODUCTION
 Materials and methods
 Results and discussion
 Conclusions
 REFERENCES
 




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C. Liu, P. M. Huang, and J. M. Zhou
Residence Time Effect on Iron Perturbation of Taranakite Formation
Soil Sci. Soc. Am. J., January 1, 2002; 66(1): 109 - 116.
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