Effect of pH on the Reductive Dissolution Rates of Iron(III) Hydroxide

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Langmuir 1997, 13, 1835-1839

1835

Effect of pH on the Reductive Dissolution Rates of Iron(III) Hydroxide by Ascorbate Yiwei Deng Department of Chemistry, Florida International University, Miami, Florida 33199 Received July 16, 1996. In Final Form: December 31, 1996X The kinetics of reductive dissolution of iron(III) hydroxide by ascorbate have been investigated as a function of pH and temperature. The reductive dissolution rate is relatively constant in the pH range of 4.3-6.0 and decreases markedly from pH 6.0 to 7.6. This pH effect is examined (1) by evaluating the dependence of the reductive dissolution rate of iron(III) hydroxide on the extent of ascorbate adsorption and on temperature and (2) by investigating the adsorption of iron(II) onto the surface of iron(III) hydroxide as a function of pH in the absence and in the presence of ascorbate. The experimental results demonstrate that the rate of reductive dissolution of iron(III) hydroxide is proportional to the extent of ascorbate adsorption over the pH range of 4.3-7.6. The dissolution rate shows the same pH-dependence trend as the surface coverage of ascorbate. The adsorption of iron(II) onto the surface of iron(III) hydroxide occurs above pH 7.0 in the absence of ascorbate but is pH independent in the presence of ascorbate. Thus, the pH effect on the reductive dissolution rate of iron(III) hydroxide can be interpreted by the pH-dependence of ascorbate adsorption. Using the data of dissolution rate and temperature, the apparent activation energies (Ea) of the overall dissolution reaction are estimated as 27 kJ/mol at pH 6.0 and 47 kJ/mol at pH 7.5. The high values of the activation energy suggest that the overall dissolution process is controlled by the surface reaction. A higher activation energy barrier must be overcome when the dissolution reaction shifts from pH 6.0 to 7.5.

Introduction Precipitation and dissolution of iron(III) hydroxides are two major processes in the redox cycling of iron in natural waters.1-5 Field investigations have shown that a variety of iron(III)-rich particles are accumulated at the redox boundary of natural waters and characterized mostly as amorphous iron(III) hydroxides.3,5,6 The iron(III)-rich particles may be produced by oxygenation of iron(II) in the oxic zone and subsequently settle into the anoxic zone of natural waters, where they can be reductively dissolved to iron(II) in the presence of reductants (e.g., products of the decomposition of biological materials,5,7 H2S(aq)8 ) and in the presence of microorganisms.9,10 Many laboratory studies have been devoted to understanding the redox cycling of iron in natural waters, particularly the mechanisms for the reductive dissolution of iron(III) hydroxides by organic reductants.11-16 It is frequently observed that the reductive dissolution of ironX Abstract published in Advance ACS Abstracts, February 15, 1997.

(1) Davison, W. In Chemical Processes in Lakes; Stumm, W., Ed.; Wiley-Interscience: New York, 1985; Chapter 2, p 31. (2) Davison, W. Earth-Sci. Rev. 1993, 34, 119. (3) Buffle J.; De Vitre R. R.; Perret D.; Leppard G. G.Geochim. Cosmochim. Acta 1988, 53, 399. (4) Buffle J.; van Leeuwen H. P. In Environmental Particles; Lewis Publisher: Boca Raton, FL, 1993, Vol. 1. (5) Balistrieri, L. S.; Murry J. W.; Paul B. Limnol. Oceanogr. 1992, 37, 510. (6) Perret D.; De Vitre R. R.; Leppard G. G.; Buffle J. In Large Lakes; Tilzer, M. M., Serruya C., Eds.; Springer-Verlag: Berlin and Heiderberg, 1990; p 224. (7) Jauregui, M. A.; Reisenauer, H. M. Soil Sci. Soc. Am. J. 1982, 46, 314. (8) Thamdrup, B.; Fossing, H.; Jorgensen B. B. Geochim. Cosmochim. Acta 1994, 58, 5115. (9) Lovley, D. R.; Phillips, E. J. P. Environ. Sci. Technol. 1991, 25, 1062. (10) Roden, E. E.; Zachara, J. M. Environ. Sci. Technol. 1996, 30, 1618. (11) Suter D.; Banwart S.; Stumm W. Langmuir 1991, 7, 809. (12) Lakind, J. S.; Stone, A. T. Geochim. Cosmochim. Acta 1989, 53, 961. (13) Dos Santos Afonso, M.; Morando, P. J.; Blesa, M. A.; Banwart, S.; Stumm, W. J. Colloid Interface Sci. 1990, 138, 74. (14) Baumgartner, E.; Blesa, M. A.; Maroto, A. J. G. J. Chem. Soc., Dalton Trans. 1982, 1649.

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(III) hydroxides by ascorbate, phenols, thiol-containing compounds, and fulvic and humic acids occurs significantly under acidic condition (pH < 5).11-16 The rate of reductive dissolution of iron(III) hydroxides is proportional to the extent of the reductant adsorption at given pH value and decreases dramatically with increasing pH.11-15 To elucidate this pH effect, several explanations have been proposed: (1) The rate of reductive dissolution is proportional to the quantity of the reductant adsorption which depends on pH;11,13 (2) the readsorption of iron(II) onto the surface of iron(III) hydroxides above pH 5 may block the surface sites and thus limit the further dissolution.12 In view of previous studies on the reductive dissolution of iron(III) hydroxides, the dissolution rates were usually derived from linear regression of the dissolved iron(II) as a function of time.11-13 Few data for the reductive dissolution rate of iron(III) hydroxides above pH 5 were reported because the concentration of dissolved iron(II) over the course of the dissolution remained below the detection limit.11-13 The pH effect could not be critically evaluated without knowing the kinetics and behavior of the reductive dissolution of iron(III) hydroxides occurring in the neutral pH range. Furthermore, the lack of direct evidence for the readsorption of iron(II) on the surface of iron(III) hydroxides in the presence of a reductant cast some doubt on the validity of the second explanation for the pH effect given above. Beside pH, temperature is recognized as another important factor that influences the dissolution rate of minerals.17,18 The temperature dependence of the dissolution rate can be generally described by the Arrhenius equation:

rate ) Ae-Ea/RT

(1)

where A is the pre-exponential factor, Ea is the activation (15) Voelker-Bartschat, B. M. Ph.D. Thesis, No. 10901, Swiss Federal Institute of Technology (ETH Zu¨rich), 1994. (16) Waite, T. D.; Morel, F. M. M. Environ. Sci. Technol. 1984, 18, 860. (17) Ganor, J.; Mogollon, J. L.; Lasaga, A. C. Geochim. Cosmochim. Acta 1995, 59, 1037. (18) Carroll, S. A.; Walther, J. V. Am. J. Sci. 1990, 290, 797.

© 1997 American Chemical Society

1836 Langmuir, Vol. 13, No. 6, 1997

energy, R is the gas constant, and T is the absolute temperature. Evaluation of the dissolution rate as a function of temperature and of solution pH allows the pH dependence of the activation energy to be determined. The values of activation energy are useful for determining the mechanisms of the overall dissolution reactions. For example, diffusion-controlled reactions have much lower activation energies (Ea < 21 kJ/mol) than the surfacecontrolled reactions (Ea ≈ 42-84 kJ/mol).19,20 The objective of this work was to investigate the pH effect on the rate of reductive dissolution of amorphous iron(III) hydroxide by ascorbate as a reductant. The iron(III) hydroxide used in this study was prepared under the conditions similar to those at the redox boundary of natural waters.21 The reactivity of iron(III) hydroxide produced, as reflected by its tendency to dissolve, was higher than that of the highly crystalline iron(III) hydroxides, which were used mostly in previous studies.21 Therefore, the rate of reductive dissolution of iron(III) hydroxide prepared in this study could be measured above pH 5 from linear regression of the dissolved iron(II) as a function of time. The pH effect on the reductive dissolution rate was then examined (1) by evaluating the dependence of the reductive dissolution rate on the extent of ascorbate adsorption and on temperature and (2) by investigating the adsorption of iron(II) species onto the surface of iron(III) hydroxide in the absence and in the presence of ascorbate over the pH range of 4.0-10. The values of the activation energy have been estimated using the Arrhenius equation. These values can provide unique insights into the reaction mechanism. Experimental Section Materials. All chemicals used in this study were analytical grade and were used without further purification. All solutions were prepared with doubly deionized water (Barnstead Nanopure). The trace amount of oxygen contained in the pure nitrogen gas (N2, 99.999%) was further removed by passing N2 through a solution of vanadium(II) with amalgamated zinc.22 Formation of Iron(III) Hydroxide. A solution (500 mL) containing 1.0 × 10-2 M NaClO4 and 3.0 × 10-4 M NaHCO3 was bubbled with pure CO2 (PCO2 ) 1 atm) until a pH value of 4.5 was attained, then 0.0183 g of Fe(ClO4)2‚6H2O was added. The initial concentration of iron(II) was 1.0 × 10-4 M. The solution was then bubbled with air (PO2 ) 0.2 atm) until pH 5 was achieved. The pH of the solution was adjusted to pH 7 by addition of imidazole buffer solution (pKa ) 7).23 The final concentration of imidazole was 0.01 M. The oxygenation of iron(II) and the subsequent precipitation of iron(III) hydroxide proceeded rapidly at about pH 7. The iron(III) hydroxide obtained was characterized by X-ray powder diffraction and by electron transmission microscopy. The iron(III) hydroxide did not show a satisfactory X-ray diffraction pattern. The electron micrograph showed iron hydroxide with a needle shape of approximately 40 × 3.4 × 2.0 nm3.24 The surface area of the iron hydroxide was estimated as 6.8 × 102 m2 g-1 (see Appendix). Reductive Dissolution of Iron(III) Hydroxide. The dissolution experiments were carried out in a 500-mL glass vessel in the absence of light (wrapped with aluminum foil). All experiments were conducted at constant temperature (20 °C) by keeping the glass vessel in a thermostated water bath. During the dissolution experiments, the pH was kept constant by addition of HClO4 (0.01 M) solution with a pH stat (Metrohm). The suspension of the freshly formed iron(III) hydroxide (1.0 × 10-2 (19) Lasaga, A. C. J. Geophys. Res. 1984, 89, 4009. (20) Locker, L. D.; de Bruyn, P. L. J. Electrochem. Soc. 1969, 116, 1659. (21) Deng, Y.; Stumm, W. Appl. Geochem. 1994, 9, 23. (22) Wersin, P. Ph.D. Thesis, No. 9230, Swiss Federal Institute of Technology (ETH Zu¨rich), 1990. (23) Martell, A. E.; Smith, R. M. In Critical stability constants: Other organic ligands; Plenum Press: New York, 1977; Vol. 3. (24) Deng, Y. Ph.D. Thesis, No. 9724, Swiss Federal Institute of Technology (ETH Zu¨rich), 1992.

Deng g L-1 as Fe(OH)3) was purged with N2 gas to eliminate the dissolved oxygen for 15 min before the required quantity of the freshly prepared ascorbate solution was added. The suspension of iron(III) hydroxide was continuously purged with N2 throughout the experiments. After addition of ascorbate solution, the aliquots of the suspension were withdrawn periodically by a syringe and filtered under nitrogen through a 0.1-µm cellulose nitrate filter (Millipore). The filtrate was collected in glass vials to which 0.1 mL of 1.0 M HNO3 had been added to quench autooxidation of iron(II). The iron(II) was analyzed using the phenanthroline method.25 The relative error and the standard deviation for measurement of iron(II) are 2.5% and 0.031, respectively. The rate of the reductive dissolution of iron(III) hydroxide can be calculated from linear regression of the dissolved iron(II) as a function of time. The overall error of the calculated dissolution rate is less than 5%. Adsorption of Ascorbate. The radioisotope 14C-labeled ascorbate (Amersham International, England) was used for the experiments of ascorbate adsorption. The stock solution of ascorbate (0.48 M) was prepared by mixing 1.0 mL of 10 µCi/mL ascorbate (4.5 × 10-4 M) with nonradioactive ascorbate (9.5 g of sodium ascorbate) and diluting to 100 mL with water. The experiments for the adsorption of ascorbate onto the surface of iron(III) hydroxide were carried out in a 25-mL flask. To initiate the adsorption experiment, a small volume (2.5 mL) of the stock solution of 14C-labeled ascorbate was added to the freshly formed iron(III) hydroxide suspension. The samples containing iron(III) hydroxide (1.0 × 10-2 g L-1 as Fe(OH)3) and ascorbate (4.8 × 10-2 M) at different pH values were equilibrated at 20 °C with shaking for 15 min (the preliminary experiment showed that the time needed to reach the adsorption equilibrium was about 8-15 min). The suspension was then filtered through a 0.1-µm filter. After filtration, the particles of iron(III) hydroxide on the filter were rinsed three times with deionized water to remove the unadsorbed ascorbate. To measure the quantity of adsorbed ascorbate, the particles of iron(III) hydroxide on the filter were dissolved with 1 mL of concentrated HCl in a sample bottle. The filter was then rinsed with 1 mL of deionized water, which was also added to the sample bottle. The contents of the sample bottle were introduced into 9 mL of xylene scintillation fluid (Lumac Gel, The Netherlands) and measured using a BETAmatic liquid scintillation counter (Kontron). The amount of adsorbed ascorbate was calculated based on the standard calibration curve. The relative error for the determination of ascorbate was not greater than 5%. Adsorption of Ferrous Iron on the Surface of Iron(III) Hydroxide. The experiments for the adsorption of iron(II) on the surface of iron(III) hydroxide were carried out in the presence and absence of ascorbate. (1) Preparation of Iron(II) Stock Solution in the Absence of Ascorbate. A solution (500 mL) containing 1.0 × 10-2 M NaClO4 and 3.0 × 10-4 M NaHCO3 was purged with pure CO2 (PCO2 ) 1 atm) until a pH value of 4.5 was achieved, and 0.0183 g of Fe(ClO4)2‚6H2O was then added. The concentration of iron(II) was 1.0 × 10-4 M. (2) Preparation of Iron(II) Stock Solution in the Presence of Ascorbate. A solution (500 mL) containing 1.0 × 10-2 M NaClO4 and 3.0 × 10-4 M NaHCO3 was purged with pure CO2 (PCO2 ) 1 atm) until a pH value of 4.5 was achieved, and 0.028 g of Fe(ClO4)2.6H2O and 9.5 g of ascorbate (sodium ascorbate) were added to the solution. The concentrations of iron(II) and ascorbate were 1.5 × 10-4 and 9.6 × 10-2 M, respectively. (3) Adsorption of Iron(II) on the Surface of Iron(III) Hydroxide. The freshly formed iron(III) hydroxide suspension (250 mL, 1.0 × 10-2 g L-1 as Fe(OH)3) was purged with pure CO2 (PCO2 ) 1 atm) until the pH reached 5.0. A 250-mL portion of the iron(II) stock solution in the absence of ascorbate was then mixed with the iron(III) hydroxide suspension at about pH 5. This gave the initial concentrations of iron(III) hydroxide of 5.0 × 10-3 g L-1 and iron(II) of 5.0 × 10-5 M, respectively. The pure CO2 gas was then substituted by nitrogen gas (N2). This resulted in a pH increase because the CO2 that had been dissolved in the suspension was driven out of the suspension. In order to obtain pH values higher than 7.5, a small volume of NaOH solution (25) Tamura, H.; Goto, K.; Yotsuyanagi, T.; Nagayama, M. Talanta 1974, 21, 318.

Effect of pH on Dissolution Rates

Figure 1. Iron(II) released as a function of time during the dissolution of iron(III) hydroxide at various pH values in the presence of ascorbate: formation of Fe(OH)3(s) (1.0 × 10-2 g L-1), imidazole 0.01 M, Fe(II)0 ) 1.0 × 10-4 M, pH 7 ( 0.1, PO2 ) 0.2 atm; dissolution of Fe(OH)3(s), ascorbate 4.8 × 10-2 M, N2, 20 °C. pH was adjusted by adding small quantities of acid, and the data refer to pH values measured. (0.01 M) was added into the suspension using a pH stat (Metrohm). The sample was taken by a syringe and filtered under nitrogen through a 0.1-µm filter. The concentration of ferrous iron was determined in the filterable fraction with the phenanthroline method.25 The same procedure was used for iron(II) adsorption onto the iron(III) hydroxide in the presence of ascorbate. Temperature Effect on the Reductive Dissolution of Iron(III) Hydroxide by Ascorbate. The experiments for the reductive dissolution of iron(III) hydroxide (1.0 × 10-2 g L-1 as Fe(OH)3) by ascorbate (4.8 × 10-2 M) at different temperatures followed the procedure for the reductive dissolution of iron(III) hydroxide described previously in this section. The dissolution experiments were carried out in a 150-mL glass vessel with a water jacket connected to the inlet and outlet of a water bath. The desired temperature was controlled by circulating the water from the water bath through the jacket of the vessel. The combined pH electrode (Metrohm) was standardized at experimental temperature against the standard buffer solutions of pH 4 and 7 (Merck).

Results and Discussion Dependence of the Reductive Dissolution Rate on pH. Figure 1 shows the concentration of dissolved ferrous iron as a function of time at various pH values upon the reductive dissolution of iron(III) hydroxide with ascorbate as a reductant. The nonzero intercepts observed in Figure 1 indicate that the dissolution rates of iron(III) hydroxides are initially rapid within 5 min over the pH range of 4.3-6.3. The rapid dissolution may be due to the dissolution of iron(III) hydroxide occurring preferentially at surface defects, such as adatom, kink, and edge sites.26 The reductive dissolution rates at different pH values are calculated from linear regression of the dissolved iron(II) as a function of time for data collected after the initial rapid dissolution. As shown in Figure 2, the dissolution rate is relatively constant in the pH range of 4.3-6.0 and decreases significantly from pH 6.0 to 7.6. Dependence of the Reductive Dissolution Rate on Ascorbate Adsorption. The extent of adsorbed ascorbate on the surface of iron(III) hydroxide was determined as various pH values (Figure 3). The ascorbate adsorption (Figure 3) exhibits the same pH-dependence trend as the rate of reductive dissolution of iron(III) hydroxide (Figure 2). In order to evaluate the dependence (26) Eggleston, R. A.; Buseck, P. R. Clays Clay Miner. 1980, 28, 173.

Langmuir, Vol. 13, No. 6, 1997 1837

Figure 2. pH dependence of the reductive dissolution rate of iron(III) hydroxide. The dissolution rates were derived from linear regression of the dissolved iron(II) as a function of time for data presented in Figure 2.

Figure 3. Surface concentration of ascorbate as a function of pH. The conditions for the formation of iron(III) hydroxide were the same as those described in Figure 1. The conditions for the adsorption of ascorbate onto the iron(III) hydroxides are as follows: Fe(OH)3 1.0 × 10-2 g L-1, ascorbate 4.8 × 10-2 M, 20 °C.

of the reductive dissolution rate on adsorption of ascorbate, the dissolution rate (log R) of iron(III) hydroxide was plotted versus the surface concentration of ascorbate (log{>Fe(III)-HA}) (Figure 4). The plot gives a satisfactory proportionality between the dissolution rate and the quantity of adsorbed ascorbate over the pH range of 4.37.6

R ) k{>Fe(III)-HA}

(2)

where k is the dissolution rate constant and {>Fe(III)HA} (mol m-2) denotes the total surface concentration of ascorbate. This proportionality indicates that the pH effect on the reductive dissolution rate of iron(III) hydroxide may be accounted for by the pH-dependence of ascorbate adsorption. A case study of the reductive dissolution of iron(III) hydroxide by Dos Santos and Stumm has shown that the reductive dissolution rate of hematite (R-Fe2O3) with H2S was proportional to the extent of surface complexes of {>Fe-S-} and/or {>FeHS} in the pH range of 5-7.27 The pH effect on the reductive dissolution rate observed in their study can be interpreted by the pH dependence of the extent of surface complex formation with sulfide. The proportionality of (27) Dos Santos, M.; Stumm, W. Langmuir 1992, 8, 1671.

1838 Langmuir, Vol. 13, No. 6, 1997

Deng

Figure 6. Temperature effect on the reductive dissolution of iron(III) hydroxide: (a) dissolution reaction at pH 6.0; (b) dissolution reaction at pH 7.5. Experimental conditions for the formation and dissolution of iron(III) hydroxide were the same as those described in Figure 1. Figure 4. Reductive dissolution rate of iron(III) hydroxide as a function of the extent of adsorbed ascorbate in the pH range of 4.0-7.6. The conditions for the formation and dissolution of iron(III) hydroxides were the same as the those described in Figure 1. The conditions for the adsorption of ascorbate onto the surface of iron(III) hydroxide were the same as those described in Figure 3.

the dissolution rate of iron(III) hydroxide to the extent of surface complex formation has also been observed in the nonreductive dissolution of lepidocrocite (γ-FeOOH) with EDTA (ethylenediaminetetraacetate) in the pH range of 3-11.28 Therefore, this inference may represent a generalized information on both reductive and nonreductive dissolution of iron hydroxides. Readsorption of Ferrous Iron on the Surface of Iron(III) Hydroxide. One of the interpretations for the pH effect on the reductive dissolution rate is that the iron(II) readsorbs on the surface sites of the iron(III) hydroxide above pH 5, thus limiting the further dissolution reaction.12,29 In order to see whether the dissolved iron(II) can be readsorbed onto the surface of iron(III) hydroxide, we examined the adsorption behavior of iron(II) in the absence and in the presence of ascorbate in the pH range of 4.0-10. Figure 5 shows the adsorption of ferrous iron on the surface of iron(III) hydroxide as a function of pH.

In the absence of ascorbate, the ferrous iron remained in the dissolved fraction from pH 4.5 to 7.2 and diminished markedly from pH 7.2 to 8.0. In the presence of ascorbate, however, almost all of the iron(II) species remained in the solution over the pH range of 5.5-10. This observation indicates that ascorbate prevents iron(II) from readsorbing onto the surface of iron(III) hydroxide above pH 7.0. Two factors may be considered to account for this observation. First, a large amount of ascorbate (4.8 × 10-2 M) competes with the iron(II) ions and/or iron(II)-ascorbate species (the stability constant for the iron(II)-ascorbate complex [Fe(II)-HA]: log K ) 1.99)22 for the surface functional groups of the iron(III) hydroxide. Second, the trace amount of oxygen, which could lead to the fast oxidation of iron(II) and the subsequent precipitation of iron(III) hydroxide above pH 5, may be removed through the redox reaction with ascorbate. Thermodynamically, the oxidation of ascorbate by oxygen (the standard redox potential E°H for the O2/H2O couple is 1.23 V)30 may occur preferentially because the standard redox potential (E° ) 0.39 V)31 for the ascorbate redox couple is lower than that (E° ) 1.0 V)20 for the [iron(III) hydroxide]/[iron(II)] redox couple. As a result, the iron(II) can be prevented from oxygenation and thus remains in liquid phase. This experimental result demonstrates that the pH effect on the rate of reductive dissolution observed here is not due to the readsorption of iron(II) species onto the surface of iron hydroxide above pH 5, which limits the further dissolution. Temperature Effect on the Reductive Dissolution of Iron(III) Hydroxide. Figure 6 shows the temperature effect on the reductive dissolution of iron(III) hydroxide with ascorbate (4.8 × 10-2 M) at pH 6.0 and 7.5. The dissolution rates of iron(III) hydroxide at different temperatures are calculated from linear regression of the dissolved iron(II) as a function of time for the data collected after the rapid dissolution. Plotting the dissolution rate against 1/T (Figure 7), one can calculate the activation energies using the Arrhenius equation (eq 1). The values of the activation energies are 27 kJ/mol at pH 6 and 47 kJ/mol at pH 7.5. The high values of activation energy are characteristics of a surface-controlled reaction rather than a transport-controlled reaction, for which typical activation energies are less than 21 kJ/mol.19,20 It is important to note that interpretation of the Arrhenius parameters based on the activated complex theory can only be applied to an elementary reaction. The

(28) Bondietti, J.; Sinniger, J.; Stumm, W. Colloids Surf. A 1993, 79, 157. (29) Mulvaney, P.; Cooper, R.; Grieser, F.; Meisel, D. Langmuir 1988, 4, 1206.

(30) Stumm, W.; Morgan, J. J. Aquatic Chemistry; Wiley Interscience: New York, 1981; p 461. (31) Clark, W. M. Oxidation-reduction potentials of organic system; The Williams & Wilkins Co.: Baltimore, MD, 1960; pp 468-470.

Figure 5. pH effect on the adsorption of iron(II) at the surface of iron(III) hydroxide (5.0 × 10-3 g L-1 as Fe(OH)3(s)): (a) adsorption of iron(II) in the absence of ascorbate (Fe(II) 5.0 × 10-5 M, N2, 20 °C); (b) adsorption of iron(II) in the presence of ascorbate (ascorbate 4.8 × 10-2 M, Fe(II) 8.0 × 10-5 M, N2, 20 °C).

Effect of pH on Dissolution Rates

Langmuir, Vol. 13, No. 6, 1997 1839

Acknowledgment. I thank Werner Stumm (Swiss Federal Institute of Environmental Sciences and Technology (EAWAG)) for simulating discussion and helpful advice throughout this study. I am greatly indebted to Rudolf Giovanoli, University of Bern, for the characterization of iron(III) hydroxide using X-ray diffraction and electron microscopy. I thank two reviewers for their valuable comments and suggestions. This work was supported by EAWAG and the Swiss National Science Foundation. Appendix: Calculation for the Surface Area of Iron Hydroxide Formed in Imidazole Buffer System

Figure 7. Arrhenius plot of the reductive dissolution rate of iron(III) hydroxide with ascorbate. The dissolution rates were calculated from linear regression of the dissolved iron(II) as a function of time for the data presented in Figure 6. The values of pH and temperature taken from Figure 6 were measured experimentally.

activation energy of the overall dissolution is considered as the apparent activation energy because the dissolution process is a complex kinetic system. It consists of several elementary steps, including diffusion of ascorbate to the surface of iron(III) hydroxide, surface adsorption and complexation, redox reaction between the adsorbed ascorbate and iron(III) center on the surface, producing iron(II) on the surface, and detachment of iron(II) from the surface into solution.11,32 It is generally believed that the attachment of reductants to the surface site is fast.12,32 The subsequent detachment of metal species from the surface is slow and requires a considerable activation energy.19,32 Therefore, the detachment step may likely control the overall rate of dissolution reaction. Conclusions The pH effect on the reductive dissolution rate of iron(III) hydroxide with ascorbate can be generally interpreted by the pH dependence of the extent of ascorbate adsorption. This finding is consistent with previous studies by Stumm and his co-workers11,27,28 and may represent a generalized interpretation for the pH effect on the dissolution of iron(III) hydroxides observed in both this study and the previous investigations. The experimental result for the adsorption of iron(II) onto the surface of iron(III) hydroxide in the presence of ascorbate demonstrates that the pH effect on the reductive dissolution rate observed in this study is not due to the readsorption of iron(II) species onto the surface above pH 5, which limits the further dissolution. The apparent activation energy (Ea) of the overall dissolution of iron(III) hydroxide was evaluated using the Arrhenius equation. It increases from 27 kJ/mol at pH 6 to 47 kJ/mol at pH 7.5. Thus, a higher energy barrier must be overcome when the overall dissolution reaction shifts from pH 6.0 to 7.5. The high activation energy values suggest that the overall dissolution process is controlled by the surface reaction. (32) Stumm, W.; Wieland, E. In Aquatic Chemical Kinetics; Stumm, W., Ed.; Wiley-Interscience: New York, 1990; Chapter 13, p 367.

The geometry of the iron hydroxide particles formed in imidazole buffer system was estimated as 40 nm × 3.4 nm × 2.0 nm according to the electron micrograph.24 The density (D) of an amorphous iron hydroxide is 2.4 g cm-3 taken from the Handbook of Chemistry and Physics.33 On the basis of the above information, the surface area of iron(III) hydroxide formed in imidazole system can be calculated as follows: For a single iron(III) particle:

surface area S ) {[2(40 × 3.4) + 2(2.0 × 40) + 2(3.4 × 2.0)] nm2} (1.0 × 10-14 cm2/nm2) ) 4.46 × 10 -12 cm2 volume

V ) [(40 × 3.4 × 2.0) nm3](1.0 × 10-21 cm3/nm3) ) 2.72 × 10-19 cm3

mass W ) VD ) 2.72 × 10-19 cm3 × 2.4 g cm-3 ) 6.53 × 10-19 g Therefore, the surface area St (in m2 g-1) of the iron hydroxide is:

St ) S/W )[4.46 × 10-12 cm2/6.53 × 10-19 g] ) 6.83 × 106 cm2 g-1 ≈ 6.80 × 102 m2 g-1 LA9607013

(33) Weast, R.; Astle, M. J.; Beyer, W. H. In Handbook of Chemistry and Physics, 66th ed.; CRC Press, Inc.: Boca Raton, FL, 1986.