Chong
Dai
a,
Juanjuan
Liu
ab and
Yandi
Hu
*a
aDepartment of Civil & Environmental Engineering, University of Houston, Houston, TX 77004, USA. E-mail: yhu11@ uh.edu; Web: http://www.cive.uh.edu/faculty/hu Fax: +(713)743 4260; Tel: +(713)743 4285
bCollege of Natural Resources and Environment, Northwest A&F University, Yangling, Shanxi 712100, China
First published on 14th November 2017
During ferrihydrite precipitation, metal ions can be sequestered in it to form impurity-bearing ferrihydrite (IBF). Using grazing-incidence small-angle X-ray scattering (GISAXS), heterogeneous precipitation/deposition of pure and IBF nanoparticles on quartz (SiO2) and corundum (Al2O3) was quantified in 0.1 mM Fe3+ solutions in the absence and presence of 1 mM Mn2+ or Al3+ (pH = 3.8 ± 0.1). The impurity ions (Mn and Al) greatly affected ferrihydrite nanoparticle precipitation/deposition on substrates. On SiO2, ferrihydrite nanoparticle precipitation/deposition was promoted in the presence of Mn but was inhibited in the presence of Al. On Al2O3, Mn- and Al-bearing ferrihydrite nanoparticle precipitation/deposition was slower than for pure ferrihydrite. Compared with on SiO2, pure and IBF nanoparticle precipitation/deposition on Al2O3 was significantly inhibited. To understand the mechanisms, interactions among impurity ions, substrates, and precipitates were explored. Surface enrichment of Mn and Al on precipitates was found to increase the zeta potential of ferrihydrite nanoparticles. The changes in surface charges of the precipitates and substrates affected heterogeneous IBF precipitation/deposition significantly. The rates and mechanisms of heterogeneous IBF precipitation/deposition provided here can help predict pollutant transport and design catalyst synthesis.
Environmental significanceThe importance of impurity-bearing ferrihydrite (IBF) nanoparticle precipitation/deposition cannot be overstated. In many natural soil and aquatic systems, IBF nanoparticle precipitation/deposition controls the fate of many aqueous contaminants. Also, IBF nanoparticles, as an economic and environmentally friendly material with large specific surface area and high reactivity, have been used widely for the removal and degradation of pollutants. To improve their performance, impurity ions are usually doped in ferrihydrite nanoparticles and the nanoparticles are usually coated on supporting substrates. Understanding the fundamentals of interactions among impurity ions, IBF nanoparticles, and substrates, which controls the heterogeneous precipitation/deposition of IBF nanoparticles, is of essential importance. |
The homogeneous precipitation of IBF has been studied widely, and impurity ions have been reported to be sequestered in ferrihydrite through surface adsorption, surface precipitation, and structural incorporation.1–4 Impurity ions have also been reported to affect the morphology, structure, and reactivity of IBF nanoparticles significantly.14,16,20–22 For example, the amount of aluminum doping in ferrihydrite can affect the sizes of the nanoparticles significantly, as well as their adsorption capacities for As(III) and As(V) and their redox reactivity for hydroquinone degradation.14,20,23,24
However, the fundamental knowledge of heterogeneous precipitation/deposition of IBF nanoparticles is much less well understood, and the controlling mechanisms are not clear.25–28 In our recent studies, different adsorption behaviors of impurity metal ions (e.g., Pb, Cu, Al and Cr) onto quartz (SiO2) were observed, which affected heterogeneous IBF precipitation/deposition on SiO2 significantly.13,29,30 In the work of Legg et al. and Tosco et al., pure ferrihydrite deposition under different ionic strength (IS) on SiO2 was measured, and the controlling mechanisms were explored.31,32 IBF precipitation/deposition on different substrates has not been reported.
Also, it is important to quantify the initial IBF precipitation (i.e., nucleation and growth) and deposition from solution. In reactive transport models, the rates of mineral precipitation are predicted as a function of the rate constant (k), solution's saturation ratio (Ω), and mineral surface area (A).33 The latter is changed by nucleation and growth, and in particular, the initial surface areas of the newly precipitated minerals are calculated based on nucleation rates. Currently, due to a lack of data on initial nucleation rates, great inaccuracy has blighted estimation of the initial surface areas of minerals, leading to inaccuracies in these reactive transport models.
Here, we aimed to fill such important information gaps. The objectives of the present study were to: (1) quantify initial heterogeneous precipitation/deposition kinetics (within 30 min) of Al- and Mn-bearing ferrihydrite on SiO2 and corundum (Al2O3) using grazing incidence small angle X-ray scattering (GISAXS) under acidic conditions (pH = 3.8 ± 0.1); (2) understand fundamental mechanisms by exploring the atomic-level interactions among impurity ions, substrates, and IBF nanoparticles with an integration of interfacial characterization methods and aquatic chemistry measurements.
All chemicals (Fe(NO3)3·9H2O, Al(NO3)3·9H2O, Mn(NO3)3·4H2O, and NaNO3) were purchased from Sigma-Aldrich (Saint Louis, MO, USA) and ultrapure water (conductivity <18 mΩ) was used to prepare all solutions (Table 1). Impure ferrihydrite precipitation experiments were conducted with 0.1 mM Fe(III) and 1 mM Al(III) or Mn(II), and the solutions were labeled as “FeAl” or “FeMn”, respectively. Pure ferrihydrite precipitation experiments were also conducted with 0.1 mM Fe(III) and the solution was labelled as “Fe only”. The IS of all solutions was adjusted to be the same, with 5.9 and 2.9 mM NaNO3 added into Fe only and FeMn solutions, respectively. Using Geochemist's Workbench (GWB; Release 9.0; Aqueous Solutions, Champaign, IL, USA), the pH, IS, and saturation indices (SI) of solutions with respect to Fe(OH)3 were calculated (Table 1). The FeAl and FeMn solutions were also calculated to be undersaturated with respect to gibbsite (Al(OH)3) and amorphous Mn(OH)2.
Sample name | Mn2+, mM | Al3+, mM | Fe3+, mM | NaNO3, mM | IS,a mM | pHb | SIc |
---|---|---|---|---|---|---|---|
Note:a IS: Ionic strength.b pH: GWB calculated pH values, which were consistent with measured pH (3.8 ± 0.1).c SI: Saturation indices with respect to Fe(OH)3. SI = log(Q/Ksp), where Q is the actual dissolved composition, and Ksp(Fe(OH)3) = 10–7.22 at 20 °C was used for SI calculations based on GWB database. | |||||||
Fe only | 0 | 0 | 0.1 | 5.9 | 6.2 | 3.8 | 0.36c |
FeMn | 1 | 0 | 0.1 | 2.9 | 6.2 | 3.8 | 0.36c |
FeAl | 0 | 1 | 0.1 | 0 | 6.2 | 3.7 | 0.23c |
After alignment, 0.7 mL of a freshly prepared solution (Fe only, FeMn, or FeAl, Table 1) was injected immediately into the cell and in situ GISAXS measurement started. The deposition and precipitation of IBF nanoparticles on substrates occurred simultaneously immediately after the reaction had started. During the measurements, GISAXS scattering signals caused by particles precipitating and depositing on substrates were collected every 1 min for 30 min. At the end of each 30 min in situ GISAXS experiment, the cell was moved horizontally for 1 mm and a GISAXS image collected at a fresh spot. The GISAXS image collected at the fresh spot was similar to the last image of each in situ GISAXS measurement, indicating no X-ray damage during our measurements. GISAXS experiments were done at Beamline 12 ID-B, Advanced Photon Source at the Argonne National Laboratory (Argonne, IL, USA).
All data analyses were conducted using IRENA, GISAXS SHOP macro, and Igor Pro 6.34 (WaveMetrics, Lake Oswego, OR, USA). Background subtraction was undertaken using the first image of each in situ GISAXS experiment, where no discernible particle scattering was detected. In two-dimensional (2D) GISAXS images, no obvious scattering patterns were shown along vertical directions, indicating that the particles on substrates were not well ordered. Therefore, for simplicity, we assumed the shape of the disordered particles to be a low-resolution and highly symmetric shape, such as a sphere. The subtracted 2D scattering images were then deducted to 1D scattering intensity curves by cutting the long Yoneda wing, where the X-ray scattering signal is the strongest because of the Vineyard effect.39,40 The deducted 1D GISAXS scattering curves were plotted as scattering intensity (I) vs. scattering vector (q, with units of Å−1, reciprocally related to size) (Fig. 2).
To obtain size information of the particles on substrates, a lognormal model of non-interacting spherical particles was used to fit all GISAXS scattering curves. Good fittings were obtained (the black lines in Fig. 2 are the fitted curves, and the colored dots are the measured data), and the size evolutions of particles (as radius, R, in nm) forming on SiO2 and Al2O3 are plotted in Fig. 3A1 and A2, respectively. Lorentz-corrected intensity curves (i.e., I × q2vs. q) were also plotted (Fig. S1 in ESI†), and the total particle volumes (V, in relative units) on SiO2 and Al2O3 were calculated as the so-called invariant , as shown in Fig. 3B1 and B2, respectively.41 By assuming a spherical shape for the particles on substrates, using the total particle volume (V) and the mean particle radius (R), the numbers of particles on substrates (N) could be calculated using N = V/R.336
Fig. 1 HRXRD patterns of the precipitates formed in 0.1 mM Fe3+ solutions in the absence (Fe only) and presence of 1 mM Mn2+ (FeMn) and 1 mM Al3+ (FeAl). |
The chemical compositions of precipitates in solutions were investigated by strong acid soak (pH = 0.5 HNO3) and consecutive dilute acid wash (pH = 3 HNO3) experiments.29,30,42,43 Solutions (500 mL) (Table 1) were freshly prepared and let to stand for 30 min. Then, precipitates were collected on the centrifugal filters (Millipore). To obtain the total amounts of impurity ions in the precipitates, particles on centrifugal filters were fully dissolved by soaking in 5 mL of 2% nitric acid (pH = 0.5) overnight. The concentrations of impurity ions (Al and Mn) and Fe were measured by inductively coupled plasma-mass spectrometry (ICP-MS), and the total atomic ratios of Mn/Al over Fe (Rt,M/Fe) in the precipitates calculated.
Studies have reported that impurity ions can be sequestered in ferrihydrite through structural incorporation and surface enrichment.1–3 To explore the sequestration mechanisms of Al or Mn in the precipitates, consecutive dilute acid wash experiments (pH = 3, HNO3) were undertaken. Initially, 5 mL of dilute HNO3 (pH = 3) was used to soak the collected particles on the filter for 10 min, and to dissolve the surface layer of the precipitates. Based on previous studies, dilute acid can desorb the metal ions adsorbed on ferrihydrite surfaces without significant dissolving ferrihydrite particles (<1%) within 30 min, which was also observed here.43 The concentrations of dissolved impurity ions (Mn/Al) and Fe ions were measured by ICP-MS, and the surface atomic ratios of Mn or Al over Fe (Rs,M/Fe) were calculated. Then, consecutively, 5 mL of pH = 3 HNO3 solutions were added to soak the remaining particles for a longer duration (i.e., 30, 60, and 360 min), until the measured atomic ratios of Mn or Al over Fe did not change significantly, which represented the lattice atomic ratios of Mn or Al over Fe (Rl,M/Fe). For the heterogeneous precipitates formed on SiO2 and Al2O3 surfaces, their amounts were too few for phase identification and chemical composition analyses.
Sample name | M(OH)xn−x,a, % | R s,bM/Fe | Rl,bM/Fe | Rt,bM/Fe | ζ,c mV | ζ_SiO2,d mV | ζ_Al2O3,d mV |
---|---|---|---|---|---|---|---|
Note:a M(OH)xn−x, %: percentages of metal ions in hydrolyzed states (e.g. Al(OH)2+, Al(OH)2+, Al(OH)3), which were calculated using GWB under our experimental conditions.b Rs,M/Fe, Rl,M/Fe, and Rt,M/Fe: surface, lattice, and total atomic ratios of impurity ions (Mn or Al) over Fe in the precipitates.c ζ: zeta potentials of pure or impurity-bearing ferrhydrite nanoparticles.d ζ_SiO2 and ζ_Al2O3: zeta potentials of quartz and corundum powders in the presence of 1 mM Mn2+ or Al3+ (pH adjusted to 3.8 ± 0.1), respectively. | |||||||
Fe only | N/A | N/A | N/A | N/A | 29.1 ± 5.7 | −22.4 ± 2.6 | 20.5 ± 2.8 |
FeMn | 0.5 | 4.14 ± 1.81 | 0.0003 ± 0.0001 | 0.04 ± 0.02 | 39.9 ± 2.7 | −25.1 ± 1.9 | 27.4 ± 6.8 |
FeAl | 4.4 | 4.36 ± 0.08 | 0.03 ± 0.01 | 0.07 ± 0.01 | 40.6 ± 2.3 | 18.1 ± 3.9 | 34.3 ± 6.4 |
Different sequestration mechanisms of impurity ions in ferrihydrite have been reported in experimental and modeling studies.44–47 Mn and Al ions can substitute Fe(III) ions in the lattice and be sequestered in Fe(III) (hydr)oxides as structural incorporation.1–3,44–46 Also, Mn and Al ions can be sequestered through surface adsorption and surface precipitation, resulting in their enrichment on ferrihydrite surfaces.1–3,46,47 Here, the sequestration mechanisms of Mn and Al in IBF were investigated by comparing the surface (Rs,M/Fe) and lattice (Rl,M/Fe) atomic ratios of the metal ions over Fe. If Rs,M/Fe ≤ Rl,M/Fe, then only structural incorporation occurred. Conversely, if Rs,M/Fe > Rl,M/Fe, then surface enrichment also occurred. As shown in Table 2, Rs,M/Fe (Rs,Mn/Fe = 4.14 ± 1.81:1, Rs,Al/Fe = 4.36 ± 0.08:1) were much higher than Rl,M/Fe (Rl,Mn/Fe = 0.0003 ± 0.0001:1, Rl,Al/Fe = 0.03 ± 0.01:1) for FeMn and FeAl, indicating that Mn and Al were enriched on the surfaces of the precipitates. This is because Fe (hydr)oxides have high affinity to various metal ions, including Al and Mn.47,48
Lattice Rl,M/Fe ratios, which are indicative of structurally incorporated impurity ions in ferrihydrite, showed a ∼100-fold difference for the precipitates formed in FeAl (Rl,Al/Fe = 0.03 ± 0.01:1) and FeMn (Rl,Mn/Fe = 0.0003 ± 0.0001:1) solutions. To understand the different incorporation behaviors of the impurity ions in ferrihydrite, we first considered the size of these impurity ions. The ionic radii of impurity ions have been reported widely to affect their structural incorporation behaviors into Fe (hydr)oxide nanoparticles at neutral pH conditions. If the ionic radius of the impurity ion is similar to that of Fe3+, it is easier for the impurity ion to substitute for the Fe3+ and cause less stress in the lattice structure of Fe (hydr)oxide nanoparticles, resulting in easier structural incorporation.49 The radii of Al3+, Mn2+, and Fe3+ have been reported to be 0.0535, 0.0645, and 0.0645 nm, respectively.50–53 If the ionic radius was a major controlling factor, then Rl,Mn/Fe should be much higher than Rl,Al/Fe. However, our measurements showed the opposite result: Rl,Al/Fe was ∼100-fold higher than Rl,Mn/Fe. With regard to the effects of ferrihydrite precipitation rates on incorporation of impurity ions, Cismasu et al. synthesized two series of Al-bearing ferrihydrite particles under different precipitation rates, and the chemical compositions of Al-bearing ferrihydrite were similar.27 For Mn-bearing ferrihydrite, the effects of precipitation rates on Mn incorporation were not reported.
Then, we recalled the pathways of ferrihydrite formation. First, Fe3+ were hydrolyzed to form Fe(OH)3 monomers (Fe3+ + 3H2O → Fe(OH)3 + 3H+). Then, polymerization occurred through continuous olation (2[Fe(H2O)5OH2+] → [(H2O)4Fe-(OH)2-Fe(H2O)4]4+ + 2H2O) and oxolation (2[Fe(H2O)5OH2+] → [(H2O)4Fe–O–Fe(H2O)5]4+ + 2H2O).54–56 In our previous study, ferrihydrite precipitation experiments were also conducted at pH = 3.7 ± 0.2 with aqueous Al/Fe ratios of 1 and 5, and the total Al/Fe ratios in the precipitates were <1%, much lower than the aqueous Al/Fe ratios.13 Under acidic pH conditions (pH = 3.8 ± 0.1) in our previous and current study, ∼95% Al ions were present as unhydrolyzed Al3+.13 Considering the hydrolysis and polymerization reactions during ferrihydrite formation, we hypothesized that only hydrolyzed metal ions (e.g., Al(OH)n3−n) could be incorporated in the lattice of ferrihydrite through oxolation and olation reactions with Fe(OH)3 monomers and polymers. Hydrolysis reactions for As metal were instantaneous, much faster than those for ferrihydrite precipitation. Therefore, the hydrolyzed ion concentrations at equilibrium conditions, calculated by GWB, were used to explain their structural incorporation in ferrihydrite. Under our experimental pH condition (3.8 ± 0.1), 4.4% of Al ions were in their hydrolyzed states (Al(OH)2+, Al(OH)2+, Al2(OH)24+, Al3(OH)45+, Al(OH)4−, and Al(OH)3), whereas only 0.5% of Mn ions were hydrolyzed. The lower amount of hydrolyzed Mn ions resulted in the small percentage of lattice Mn/Fe ratio in the precipitates, a result which supported our hypothesis. This hypothesis is also supported by studies of Al incorporation during ferrihydrite precipitation at neutral pH conditions.57,58 For example, Cismasu et al. synthesized a series of Al-bearing ferrihydrite particles in solutions (pH = 7.5 ± 0.2) with a total metal ion concentration of 2 mM and various aqueous Al/Fe ratios (0.10–0.67).58 Under the neutral pH condition, as >99% of Fe and Al ions were hydrolyzed, the atomic ratios of Al/Fe in the precipitates were similar to their aqueous ratios.58
Our hypothesis was further supported by the reported Cr incorporation in ferrihydrite during ferrihydrite precipitation at neutral and acidic pH conditions.2,59,60 For example, in the work of Tang et al., a series of Cr-bearing ferrihydrite nanoparticles was synthesized with a total metal ion concentration of 0.1 M and different Cr/Fe ratios (0.1–9) at pH ∼ 7.2 Under this neutral pH condition, because >99% Fe and Cr ions were hydrolyzed, the atomic Cr/Fe ratios in the Cr-bearing ferrihydrite particles were similar to the initial aqueous Cr/Fe ratios.2 Conversely, under an acidic condition (pH = 3.7 ± 0.1), because only ∼36% of aqueous Cr ions were hydrolyzed, the atomic ratios of Cr/Fe (0.06–0.18) in the precipitates were much lower than the aqueous Cr/Fe ratios (0.2–2.5).30
To summarize, all previous studies of Al and Cr incorporation in ferrihydrite under acidic and neutral pH conditions confirmed our hypothesis: i.e., the hydrolysis of impurity ions affected their structural incorporation in ferrihydrite.2,13,30,57–60 IBF precipitation experiments have been conducted,2,13,30,57–60 and the ionic radii of metal ions were considered to be the controlling factors for their structural incorporation during IBF precipitation. Here, the hydrolysis of impurity ions, which is highly dependent upon pH, was reported, for the first time, to affect their structural incorporation in IBF nanoparticles greatly. The fundamental information provided here can be used to better understand and control incorporation of impurity ions in ferrihydrite during its precipitation under different pH conditions. Also, it would be an interesting future direction to develop a correlation equation of incorporation of impurity ions in ferrihydrite as a function of the radius, oxidation states, and hydrolysis reactions of impurity ions.
To understand the different effects of impurity ions on heterogeneous ferrihydrite precipitation/deposition, the electrostatic interactions between the precipitates and substrates were considered. The zeta potential of Mn- and Al-bearing ferrihydrite nanoparticles was measured to be 39.9 ± 2.7 and 40.6 ± 2.3 mV, higher than that of pure ferrihydrite (29.1 ± 5.7 mV). Impurity ions, adsorbed or precipitated on ferrihydrite surfaces, have been reported to change the pHpzc of ferrihydrite nanoparticles.58 With a pHpzc of 8.5 for pure ferrihydrite, the pHpzc of Si-bearing ferrihydrite nanoparticles decreased due to the low pHpzc of silica (2–4) on ferrihydrite surfaces. In contrast, the pHpzc of Mn(OH)2 and Al(OH)3 was higher (∼10) than that of ferrihydrite.61,62 Therefore, the Mn and Al enriched on ferrihydrite surfaces resulted in the increased zeta potential of Mn- and Al-bearing ferrihydrite nanoparticles.
The zeta potential of SiO2 powder in the absence and presence of NaNO3 was similar, indicating no adsorption of Na+ onto SiO2 powder under our experimental conditions (pH = 3.8 ± 0.2). The SiO2 surfaces in the absence and presence of 1 mM Mn(II) ions under our experimental pH conditions (3.8 ± 0.1) were both negatively charged, with a zeta potential of −22.4 ± 2.6 and −25.1 ± 1.9 mV, respectively. Conversely, in the presence of 1 mM Al(III) ions, the SiO2 surfaces became positively charged (zeta potential = 18.1 ± 3.9 mV). The different effects of impurity ions on the surface charges of SiO2 may be related to the different behaviors of metal ion adsorption onto substrates, which is an important topic that we are investigating now, but it is not the focus of the present study. The surface zeta potential of the precipitates and substrates, which determined their electrostatic interactions, explained well the observed trend of heterogeneous ferrihydrite precipitation/deposition. In the presence of Mn, stronger attractive forces existed between the more positively charged Mn-bearing ferrihydrite (39.9 ± 2.7 mV) and negatively charged SiO2 surfaces (−25.1 ± 1.9 mV), resulting in the promoted Mn-bearing ferrihydrite precipitation/deposition on SiO2. Conversely, in the presence of Al, repulsive forces existed between the positively charged Al-bearing ferrihydrite precipitates (40.6 ± 2.3 mV) and SiO2 surfaces (18.1 ± 3.9 mV), thereby hindering Al-bearing ferrihydrite precipitation/deposition on SiO2. The changes in zeta potential of the substrates in FeMn and FeAl solutions could be caused by adsorption of Mn2+ and Al3+ or heterogeneous precipitation of Mn(OH)2 and Al(OH)3 on substrates. Further discussion about possible adsorption of Mn2+ and Al3+vs. Mn(OH)2 and Al(OH)3 precipitation on substrates can be found in ESI.†
Heterogeneous nucleation and deposition can increase the total particle number on substrates. The deposition behavior can be predicted by the Derjaguin, Landau, Vervey, and Overbeek (DLVO) theory. For a Fe-only solution, stronger electrostatic repulsive forces were present between more positively charged Al2O3 surfaces and ferrihydrite surfaces, which should have inhibited ferrihydrite deposition onto Al2O3 to a greater extent than on SiO2. Therefore, if deposition was the dominant process, the number of ferrihydrite particles (N) on Al2O3 should be less than that on SiO. However, the number of pure ferrihydrite particles on Al2O3 was calculated to be ∼2-fold higher than that on SiO2. Therefore, heterogeneous nucleation, instead of deposition, should be the dominant pathway for the pure ferrihydrite nanoparticle formation on Al2O3.
The promoted nucleation of pure ferrihydrite on Al2O3 rather than on SiO2 has been reported in our previous work of pure ferrihydrite precipitation on SiO2 and Al2O3.36 The latter has higher hydrophobicity, which indicates higher substrate–water interfacial energy than for SiO2,36 so the interfacial energy barrier for heterogeneous ferrihydrite nucleation on Al2O3 was lower than that on SiO2 according to classic nucleation theory.36 Therefore, heterogeneous ferrihydrite nucleation was promoted on Al2O3.36
The total volume of nanoparticles (representing precipitation/deposition rates) on Al2O3 surfaces was also compared (Fig. 3B2). Compared with pure ferrihydrite precipitation on Al2O3, Al- and Mn-bearing ferrihydrite precipitation/deposition was inhibited. Such inhibition can be explained well from the zeta-potential differences of the substrates and nanoparticles in the absence and presence of impurity ions. In the presence of Al and Mn, the zeta potential of Al2O3 and precipitate surfaces was more positively charged, and the stronger electrostatic repulsive forces between the substrate and IBF nanoparticles resulted in slower IBF precipitation/deposition on Al2O3. In the presence of Al and Mn, the increase in the zeta potential of IBF nanoparticles was caused by Al/Mn enrichment on the particle surfaces, as explained above. The increase in the zeta potential of Al2O3 might have been caused by adsorption of Mn and Al ions onto Al2O3. Further understanding of the mechanisms controlling ion adsorption onto substrates is not the focus of the present paper, but is a very interesting research topic which we are currently exploring.
In summary, the hydrolysis of impurity ions affected their structural incorporation into ferrihydrite lattices. Also, enrichment of impurity ions on IBF surfaces can affect the zeta potential of IBF nanoparticles. Furthermore, the surface charges of substrates can be altered greatly in the presence of impurity ions. Through all these three mechanisms, impurity ions can affect the heterogeneous precipitation/deposition of IBF on mineral surfaces significantly.
Footnote |
† Electronic supplementary information (ESI) available: Details of GISAXS calculations of critical angles, XPS measurements, and discussion on adsorption of Mn2+ and Al3+ ion vs. Mn(OH)2 and Al(OH)3 precipitation on substrates, as well as GISAXS Lorentz-corrected curves (Fig. S1) and XPS measurements of Mn oxidation states (Fig. S2). See DOI: 10.1039/c7en00835j |
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