Reduction of crystalline iron(III) oxyhydroxides using hydroquinone: Influence of phase and particle size
© American Institute of Physics 2005
Received: 07 January 2005
Accepted: 25 July 2005
Published: 09 September 2005
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© American Institute of Physics 2005
Received: 07 January 2005
Accepted: 25 July 2005
Published: 09 September 2005
Iron oxides and oxyhydroxides are common and important materials in the environment, and they strongly impact the biogeochemical cycle of iron and other species at the Earth's surface. These materials commonly occur as nanoparticles in the 3–10 nm size range. This paper presents quantitative results demonstrating that iron oxide reactivity is particle size dependent. The rate and extent of the reductive dissolution of iron oxyhydroxide nanoparticles by hydroquinone in batch experiments were measured as a function of particle identity, particle loading, and hydroquinone concentration. Rates were normalized to surface areas determined by both transmission electron microscopy and Braunauer-Emmett-Teller surface. Results show that surface-area-normalized rates of reductive dissolution are fastest (by as much as 100 times) in experiments using six-line ferrihydrite versus goethite. Furthermore, the surface-area-normalized rates for 4 nm ferrihydrite nanoparticles are up to 20 times faster than the rates for 6 nm ferrihydrite nanoparticles, and the surface-area-normalized rates for 5 × 64 nm goethite nanoparticles are up to two times faster than the rates for 22 × 367 nm goethite nanoparticles.
Iron oxide and oxyhydroxides, hereafter referred to as iron oxides, often occur as nanoparticles in the 3–10 nm size range and are found in diverse environments, such as in weathering rinds of iron-bearing minerals or in lakes, streams, aquifers, and acid mine drainage. [1, 2] These materials can be formed and transformed by a variety of redox and nonredox processes, which can be driven both abiotically and biotically. [3, 4] For example, Thiobacillus ferrooxidans has been shown to catalyze the oxidation of Fe(II), whereas Shewanella species have been shown to reductively dissolve Fe(III) oxides. [6–15] Furthermore, iron oxides can be abiotically formed through chemical weathering of iron-bearing minerals or when aqueous fluids containing high concentrations of dissolved Fe(II) encounter fluids rich in dissolved oxygen.  Such reactions can result in phase changes through precipitation, dissolution, and solid-state phase transformations. [6, 12, 14, 16] These are often key elements of the processes governing the transport and fate of naturally occurring and anthropogenic chemical species, such as nitroaromatics and arsenic species. [1, 2, 17–19]
Previous work has examined size-dependent reactivity of nanoscale materials in catalysis,[20, 21] adsorption, [22–24] and biotic dissolution. [8, 13] Furthermore, many researchers have attempted to systematically examine the relative reactivity of iron oxide particles (e.g., Refs. 3, 6, 7, 9, 13, and 25–31, and references contained therein). Many such studies attempt to link particle size with reactivity but lack a holistic and consistent characterization of the solid materials used. Findings are often inconsistent, and differences observed are often small and difficult to explain. Finally, recent work has addressed whether reactivity trends are consistent between abiotic and biotic experiments using iron oxide particles of different phase and size. [7, 9] Much of this work suggests a particle size effect in the reactivity of the iron oxides. Comparisons between these distinct types of work could be substantially strengthened by the use of a common set of materials, especially in the case of experiments using a broad range of redox agents and microorganisms. In such a way, the idea of universal reactivity (e.g., whether similar trends are observed for abiotic and biotic reactions involving the same particles) can be tested, and we are developing a well-characterized library of procedures, characterization results, and even some solid materials that can be used by many different research groups in order to make meaningful comparisons between varieties of experiments (including both abiotic and biotic experiments) possible.
This paper presents quantitative results demonstrating that abiotic iron oxide reactivity is particle size dependent. Specifically, this paper examines the abiotic reduction of two types of well-characterized ferrihydrite (Fe5HO8·4H2O)3 and goethite (α-FeOOH) nanoparticles using hydroquinone as the reducing agent. Hydroquinone was selected because many of the biotic mechanisms involving redox of iron oxides use quinones as electron shuttles. [11, 12, 32–36] Also, quinone functional groups have been found in natural humic substances and can be generated biotically through the degradation of lignins, among other compounds. [37, 38] A number of researchers have studied the reduction of iron oxides by quinones and results show they can effectively be used to quantitatively evaluate the reactivity of these particles. [28, 39–42]
The synthesis method was adapted from Burleson and Penn.  Using a peristaltic pump at a rate of 4.58 mL/min, 1.0 L of 0.4799 M NaHCO3 (Fisher, ACS grade) was added dropwise to a continuously stirred 1.0 L solution of 0.4000 M Fe(NO3)3·9H2O (Fisher, ACS grade). During the transfer, the solution changed from bright orange to dark brownish red with no visible precipitate. The suspension was separated into 250 mL Nalgene bottles and microwaved one at a time until boiling occurred, shaking every 40 s (most boiled after 120 s). Immediately after heating, each suspension was plunged into an ice bath until it reached ~20°C. In order to remove the counter ions present from the synthesis, the cooled suspensions were placed into dialysis bags (MWCO = 2000), which were placed in Milli-Q® H2O. The water was changed three times per day for three days. The resulting suspensions were placed in a number of weigh boats, covered, and placed in a fume hood to dry. The justification for drying the particles is that previous results have shown that the six-line ferrihydrite nanoparticles prepared using this method are not stable in aqueous suspension because they grow by oriented aggregation accompanied by phase transformation to goethite in a relatively short period of time, even at room temperature (i.e., within a few weeks).  Using a mortar and pestle, the dry, dark reddish brown particles were ground into a fine powder and stored in a glass vial.
The synthesis method was adapted from Schwertmann and Cornell.  With stirring, 20 g of solid Fe(NO3)3·9H2O was added to 2.0 L of Milli-Q H2O at 75°C. The solution temperature was maintained at 75°C for 12 min. After heating, the solution was plunged into an ice bath until it reached ~20°C (~30 min). The cooled suspension was dialyzed, dried, and ground as described earlier.
The synthesis of the nanorods began by using the 4 nm-6LF synthesis procedure (Sec. II A). After dialysis, the pH of the ferrihydrite suspension was quickly adjusted to 12 using 5 M NaOH (Fisher, ACS grade). A deep maroon suspension formed. The suspension was heated at 90°C for 24 h, after which a deep orange precipitate had settled to the bottom quarter of the bottle. The supernatant was discarded, and the remaining suspension was placed into dialysis bags, which were placed in Milli-Q H2O. The water was changed three times per day for three days. The resulting suspension was dried and ground as described above.
The synthesis method was adapted from Schwertmann and Cornell.  With stirring, 40 g of Fe(NO3)3·9H20 was added to 900 mL of Milli-Q H2O. While stirring, the pH was adjusted to 12 using 5 M NaOH. A deep maroon suspension formed. The suspension was heated in an oven at 90°C for one week, after which a yellow-orange precipitate had settled to the bottom quarter of the bottle. The supernatant was discarded, and the remaining suspension was placed into dialysis bags, which were placed in Milli-Q H2O. The water was changed three times per day for three days. The resulting suspension was dialyzed, dried, and ground as described earlier.
All preparations and reactions were performed in a catalytically maintained anaerobic environment (~3% H2 in N2, vinyl anaerobic chamber, Coy Laboratory Products Inc., O2 < 100 ppm, as measured by an oxygen/hydrogen gas analyzer). Suspensions were stirred using Teflon™-coated stir bars, and all reaction bottles were covered with aluminum foil to prevent exposure to light.
Known masses (75, 50, 25, or 12.5 mg) of particles were placed in clean, 30 mL Nalgene bottles containing 5.0 mL of 40 mM, pH 3.75 acetate buffer [prepared using glacial acetic acid (Mallinckrodt, ACS grade) and NaOH] that had been purged with N2 gas for at least 20 min. The suspensions were capped, removed from the anaerobic chamber, and sonicated for 10 min. After returning the bottles to the anaerobic chamber and stirring overnight, the appropriate volume of acetate buffer (to bring the final reaction volume to 25.0 mL) and 10 mM hydroquinone (QH2, Sigma, 99%) stock solution were added. The samples were stirred continuously throughout the experiment. At desired time intervals, 1.0 mL aliquots were removed and filtered using a 0.2 μm Pall nylon filter membrane. The concentration of p-benzoquinone (Q) at time t, [Q] t , was immediately quantified (<1 min) via high performance liquid chromatography (HPLC). Stop time was recorded as the time of filtering. The solid concentration was assumed to be constant since no settling or clumping of the particles was observed during sampling.
Blank samples containing only QH2 in acetate buffer were used to account for the spontaneous oxidation of QH2 to Q.
Samples were quantified using an Agilent Technologies 1100 Series HPLC equipped with a Zorbax® C18 Stable Bond column. The flow rate was 0.75 mL/min, and the mobile phase consisted of 65 vol % 40 mM, pH 3.75 acetate buffer, and 35 vol % acetonitrile (Pharmco, HPLC grade). The detecting wavelength was 235 nm. The injection volume was 10 μL. Using this method, the retention time of QH2 was 2.4 min, and the retention time of Q was 3.4 min. An eight-point calibration curve from 0 to 1 × 10-2 M QH2 and an eight-point calibration curve from 0 M to 1 × 10-3 M Q (Acros, 99 + %) were used.
Aqueous [Fe(II)] was determined using an adaptation of the Ferrozine assay for a subset of experiments in order to confirm that the reaction proceeded via reductive dissolution and to verify the stoichiometry of the reaction for each particle type. Immediately after HPLC analysis, 0.80 mL of the filtered sample was combined with 0.25 mL of 5 g/L Ferrozine (Acros, 99.9%) and 3.25 mL 40 mM, pH 3.75 acetate buffer. Absorbance at 562 nm was measured using a Spectronic 20D+ visible spectrophotometer. A standard Ferrozine/acetate buffer solution was used as a blank. An eight-point calibration curve from 0 M to 1 × 10-4 M FeCl2·4H2O (Fisher, ACS grade) was used. Blanks containing particles in buffer (no reductant) were also tested. For at least one sample of each particle type, the amount of Fe(II) adsorption onto the nanoparticles was determined by difference between the total Fe(II) equivalents in solution and the benzoquinone equivalents produced.
Particles were characterized by three different methods: x-ray diffraction (XRD), transmission electron microscopy (TEM), and Brunauer-Emmett-Teller surface area analysis (BET). 
XRD was performed using a PANalytical X'Pert Pro theta-theta diffractometer equipped with a Co anode and an X'Celerator detector. Data collection ranges and collection times were 10° to 90° 2θ and 120 min, respectively, for ferrihydrite samples and 15° to 90° 2θ and 30 min for goethite samples. The diffraction patterns were compared to PDF (powder diffraction file) No. 29-0712 (six-line ferrihydrite) and PDF No. 29-0713 (goethite).
Surface areas (± standard deviation) for iron oxyhydroxide particles used. SABETdenotes surface area determined by the BET method, and SATEMdenotes surface area determined from particle size data obtained from TEM images (see the text).
1565 ± 27
409 ± 74
4.0 ± 0.2
234.9 ± 0.5
271 ± 49
5.9 ± 0.3
136.8 ± 0.5
210 ± 19
5.3 ± 0.3
64 ± 3
38.19 ± 0.19
53 ± 5
22 ± 1.1
367 ± 18
The anticipated half reactions and overall reaction of ferrihydrite [Eq. (1)] or goethite [Eq. (2)] with hydroquinone (QH2) can be written as
2Fe5HO8·4H2O(s) + l0e- + 30H+(aq) → 10Fe2+(aq) + 24H2O(1),
5QH2(aq) → 5Q(aq) + 10e- + 10H+(aq),
2Fe5HO8·4H2O(s) + 5QH2(aq) + 20H+(aq) → 5Q(aq) + 10Fe2+(aq) + 24H2O(l), (1)
10α- FeOOH(s) + l0e- + 30H+(aq) → 10Fe2+(aq) + 20H2O(1),
5QH2(aq) → 5Q(aq) + 10e- + 10H+(aq),
10α- FeOOH(s) + 5QH2(aq) + 20H+(aq) → 5Q(aq) + 10Fe2+(aq) + 20H2O(1). (2)
For Semiquantitative comparisons between different phases, initial rates were computed by fitting the linear portion [solid lines shown in Fig.4(a)] of the graph to a least squares regression following the method of initial rates.  Rate constants and reaction orders with respect to both QH2and the reactive surface sites can be found using Eq. (3) (adapted from the method used by Stack et al.)
d[Q]/dt = r = k[QH2] m [S] n , (3)
Reaction orders for QH2 (m) and surface (n) and rate constants (k) for semiquantitative comparison between each particle type by the method of initial rates. Errors reported are standard errors. SATEM normalized rate constant is denoted by kTEM, and SABET normalized rate constant is denoted by kBET.
k × 104 (h-1)
kTEM × 105 (h-1 m-2)
kBET× 105 (h-1m-2)
0.22 ± 0.04
1.00 ± 0.06
104 ± 6
2.6 ± 0.5
0.67 ± 0.04
0.50 ± 0.05
0.59 ± 0.06
96 ± 9
3.5 ± 0.7
4.1 ± 0.4
0.39 ± 0.09
0.13 ± 0.04
0.818 ± 0.006
0.037 ± 0.003
0.0598 ± 0.0005
0.36 ± 0.04
0.18 ± 0.07
0.21 ± 0.03
0.040 ± 0.006
0.056 ± 0.007
The method of initial rates, which uses only the initial, linear portion of the data, provides for excellent, semiquantitative comparison between the goethite and ferrihydrite samples. However, it is not adequate for a quantitative comparison between the 4 nm-6LF and 6 nm-6LF ferrihydrite samples and the nanorod and microrod goethite samples. In order to refine the quantitative comparisons, the data were fit using one-dimensional (1D) diffusion kinetics for ferrihydrite and two-dimensional (2D) diffusion kinetics for goethite. These diffusion models are based on the idea that the kinetics are governed by the movement of the particles through the solution.
In the case of the spherical ferrihydrite particles, 1D diffusion kinetics yield the best fits since the particles essentially act as spheres moving through the solution, and this is consistent with the approach used by Dold.  Using Eq. (4), in which r represents the 1D dissolution rate and α represents the fraction of solid dissolved at any given time, t, the data were fit by minimizing the unsigned mean error between the data and the model curve,
α2 = rt. (4)
Reaction orders for QH2 (m) and surface (n) and rate constants (k) for each particle type using the ID diffusion kinetic model. Errors reported are standard errors. SATEM normalized rate constant is denoted by kTEM, and SABET normalized rate constant is denoted by kBET.
k × 105 (h-1)
kTEM× 108 (h-1m-2)
kBET × 108 (h-1m-2)
0.55 ± 0.02
0.51 ± 0.06
2.02 ± 0.10
4.9 ± 0.9
1.29 ± 0.07
0.79 ± 0.10
1.24 ± 0.12
0.082 ± 0.010
0.30 ± 0.07
0.35 ± 0.04
Alternative explanations for the reactivity difference include the presence of trace goethite in 4 nm-6LF, the presence of carbonate in 4 nm-6LF, and a possible difference in crystallinity between the 4 nm-6LF and 6 nm-6LF samples. The presence of goethite in 4 nm-6LF most likely means that the reactivity difference observed is a minimum estimate because goethite particles are dramatically less reactive than ferrihydrite particles (Table II). Next, 4 nm-6LF particles were prepared using carbonate while 6 nm-6LF particles were prepared without addition of carbonate (see methods, Secs. IIA and IIB). Preliminary results using similarly sized particles that were prepared with and without carbonate suggest a small decrease in reactivity with the inclusion of carbonate (reductive dissolution is only 1.3 times faster for the carbonate-free particles). This suggests, again, that the reactivity difference observed is a minimum. Finally, it is possible that the reactivity difference is due to a crystallinity difference between the ferrihydrite samples. XRD patterns demonstrate that both ferrihydrites can be clearly classified as six-line ferrihydrite. However, the pattern for 6 nm-6LF exhibits narrower peaks than does the pattern for 4 nm-6LF, which could be attributed to both a larger particle size and a higher degree of crystallinity. TEM results demonstrate a difference in particle size, which suggests that the primary difference between these two ferrihydrites is one of size. It is likely that surface energy varies as a function of size, and such an effect could be a consequence of the relative number of atoms at kinks versus facets and the relative number of atoms contained at or near the surface in comparison to the bulk. We conclude that the size-dependent reactivity observed here is a direct consequence of a difference in surface energy between these materials and that the surface energy of 4 nm-6LF is greater than that of 6 nm-6LF. Finally, the difference in reaction orders is interesting. This result suggests that the mechanism by which the reductive dissolution occurs may change as a result of the synthetic procedure (i.e., the use of carbonate).
For the acicular goethite particles, 2D diffusion kinetics yield the best fits since the particles act as cylinders moving through the solution, and this is consistent with the approach used by Houben.  Using Eq. (5), in which r represents the 2D dissolution rate and α represents the fraction of solid dissolved at any given time, t, the data were fit by minimizing the unsigned mean error between the data and the model curve,
(1 - α)In(1 - α) + α = rt. (5)
Reaction orders for QH2 (m) and surface (n) and rate constants (k) for each particle type using the 2D diffusion kinetic model. Errors reported are standard errors. SATEM normalized rate constant is denoted by kTEM, and SABET normalized rate constant is denoted by kBET.
k × 108 (h-1)
kTEM × 1010(h-1m-2)
kBET × 1010(h-1m-2)
0.92 ± 0.12
1.61 ± 0.13
9.9 ± 1.2
4.5a ± 0.7
7.2a ± 0.9
0.82 ± 0.18
1.74 ± 0.18
1.38 ± 0.16
2.6a ± 0.4
3.6a ± 0.4
The relative reactivity of ferrihydrite and goethite nanoparticles has been quantified by measuring the rates of reductive dissolution of nanoparticles by hydroquinone. Results confirm that ferrihydrite is substantially more reactive than goethite and demonstrate that reactivity is size dependent in both goethite and ferrihydrite. While variations in synthesis methods, crystallinity, and/or presence of impurities may influence reactivity, our results show that these effects are small for this suite of samples. Experiments further exploring such effects are currently under way. Finally, these results highlight the need to develop synthetic methods that produce homogeneous and monodisperse goethite particles of a smaller size. To date, no method has been developed that can produce goethite nanoparticles in the 3–10 nm size range.
A. J. A. and R. L. P. thank the University of Minnesota and the National Science Foundation (Divisions of Chemistry and Earth Sciences) for funding. We thank S. L. Brantley, J. L. Cantolina, and J.-H. Jang at Penn State University for assistance with BET and J. E. Kabrhel for assistance with NMR.