δMnO2 is an important component of environmental minerals and is among the strongest sorbents and oxidants. The crystalline morphology of δ-MnO2 is one of the key factors affecting its reactivity. In this work, δ-MnO2 was initially synthesized and placed in an acidic environment to react with Mn2+ and undergo a crystalline transformation. During the transformation of crystalline δ-MnO2, kinetic sampling was conducted, followed by analyses of the structures and morphologies of the samples. The results showed that at pH 2.5 and 4, δ-MnO2 nanoflakes spontaneously self-assembled into nanoribbons via edge-to-edge assembly in the initial stage. Subsequently, these nanoribbons attached to each other to form primary nanorods through a face-to-face assembly along the c-axis. These primary nanorods then assembled along the (001) planes and lateral surfaces, achieving further growth and thickening. Since a lower pH is more favorable for the formation of vacancies in δ-MnO2, δ-MnO2 can rapidly adsorb Mn2+ directly onto the vacancies to form tunnel walls. At the same time, the rapid formation of the tunnel walls leads to a quick establishment of hydrogen bonding between adjacent nanoribbons, enabling the assembly of these nanoribbons into primary nanorods. Therefore, in a solution with the same concentration of Mn2+, the structure transformation and morphology evolution of δ-MnO2 to α-MnO2 occur faster at pH 2.5 than at pH 4. These findings provide insights into the mechanism for crystal growth from layer-based to tunnel-based nanorods and methods for efficient and controlled syntheses of nanomaterials.
Research Article
Effect of Mn2+ concentration on the growth of δ-MnO2 crystals under acidic conditions
https://doi.org/10.21203/rs.3.rs-4815763/v1
This work is licensed under a CC BY 4.0 License
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α-MnO2 is the predominant manganese oxide in the supergene oxidation zones of manganese-bearing crusts, manganese deposits, and lateritic weathered profiles.1–4 Representing a significant group within the transition metal oxides (TMOs), α-MnO2 features 2 × 2 and 1 × 1 tunnel structures (with occasional intergrowth of 2 × 3, 2 × 4, and “T” junction structures), which arise from double chains of edge-sharing [MnO6] octahedra corner-sharing with neighboring chains.5–7 These larger tunnels are typically stabilized by cations such as K+, Ba+, and Na+.8–12 The unique physicochemical properties of α-MnO2, including its distinct structural and morphological characteristics, enable a broad spectrum of potential applications in environmental pollution abatement, supercapacitors, and molecular sieves.13–16 The efficacy of these applications is significantly influenced by the specific structure and morphology of α-MnO2.7
Previous research has revealed that α-MnO2 forms from the transformation of δ-MnO2 in acidic environments (pH < 4).7, 17–19 δ-MnO2 is naturally formed and widely present in nature.20 It has poorly crystalline, c-axis disorder stacked [MnO6] octahedral layered structure, strong adsorption and oxidation activity.21–24 Our previous study on the one-step synthesis of α-MnO2 revealed that the initially formed δ-MnO2, driven by an increase in internal Mn(III), undergoes a multi-stage oriented assembly of δ-MnO2 nanosheets, ultimately leading to the formation of α-MnO2 nanorods.7 Grangeon et al through aging experiments of δ-MnO2 with pH 3–10, found that only the δ-MnO2 at pH 3 transformed into α-MnO2 after four years of aging.17 Above all, the coupled structural and morphological evolution from δ-MnO2 to α-MnO2 is influenced by both pH and Mn(III), although the specifics of this influence remain unclear. Understanding the morphological evolution throughout the entire synthetic process and its relationship with structural transformation, as well as the influence of pH and Mn(III) on the morphological evolution during the transformation from δ-MnO2 to α-MnO2, is critical for α-MnO2 material synthesis in various applications and for understanding the mineralogy and behavior of different types of MnO2 minerals in nature.8–16, 20, 23–27
In this study, poorly crystalline δ-MnO2 was synthesized. We investigate the influence of pH and Mn2+ on the structural transformation and morphological evolution of δ-MnO2 into α-MnO2 by introducing varying concentrations of Mn2+ at pH 2.5 and pH 4 solutions. The transformation process was observed with morphology and kinetic analyses to study the mechanism. The changes in the nanoparticle morphology and structure during the transformation of δ-MnO2 to α-MnO2 were observed by X-ray diffraction (XRD) and high-resolution transmission electron microscopy (HRTEM). This provides not only insights into the formation process of α-MnO2 nanorods but also new possibilities for controlled synthesis of α-MnO2 materials.
2.1 Experimental materials and methods
2.1.1 Reagents used in the experiment
The reagents used in this experiment included NaMnO4·H2O (> 97%), MnSO4·H2O (> 99%), and Hac (> 97%), all of which were purchased from Sigma-Aldrich (St Louis, MO). Distilled deionized water (DDW) was used in the preparations of mineral reagents and for washing the products. The DDW was obtained by treating distilled water with a Labconco Water Pro Ps system, and the resulting resistivity was greater than 18.4 MΩ•cm.
2.1.2 Synthesis of poorly crystalline δ-MnO2
In a typical experiment, 10.00 g NaMnO4 was dissolved in 680 mL of 0.25 mol/L NaOH solution. Then, 320 mL of 0.3 mol/L MnCl2 solution was added to the NaMnO4 solution dropwise at a rate of 10.00 mL/min. The suspension was vigorously stirred for an additional 2 h and then standing for 4 h. The precipitate which obtained after centrifugation was mixed with 1 M NaCl2 solution for 1 h. The suspension was filtered and washed until the conductivity of the supernatant was less than 20.0 µS cm− 1.7
2.1.3 Isothermal adsorption
The isothermal adsorption experiments were controlled variable experiments, and the controllable factors in this experiment were the concentration of Mn2+ and the pH of the solution. The different levels of each factor were set based on the relevant literature.28 The concentrations of Mn2+ were 0 mol/L, 0.05 mmol/L, 0.1 mmol/L, 0.3 mmol/L, 0.7 mmol/L, 1 mmol/L, 5 mmol/L, 10 mmol/L, 25 mmol/L, and 50 mmol/L, and the pH was adjusted to 2.5, 4, and 5. There were 3 groups of experiments (Table 1). Each set of experiments was repeated three times, and the average value was taken.
For each sample, 2.5 mL of 36 g/L suspension of poorly crystalline δ-MnO2 was added to a centrifuge tube with a certain amount of MnSO4 solution and a certain volume of deionized water. The total volume of the solution was 30 mL. The pH of the mineral suspensions was adjusted to 0.01 mol/L, 0.05 mol/L, 0.1 mol/L, 0.5 mol/L, 1.00 mol/L NaOH and H2SO4. The changes in the suspension pH were regularly monitored, the set pH was maintained, and the mixture was allowed to react for several days until the pH change was less than ± 0.05 within 7 days. After 7 days of stirring, the sample was centrifuged once in a centrifuge tube, the supernatant was removed, and the sample was washed and centrifuged several times until the conductivity of the supernatant was 20 µS/cm. After centrifugation, a high-speed freezing centrifuge (BECKMAN J2-MC) was used at relative centrifugal force (RCF) of 17800 g, and the centrifugation time was 6 minutes. The washed paste was freeze-dried, ground in an agate mortar and then bottled for later use. The concentration of Mn2+ in the supernatant after centrifugation was determined with flame atomic absorption spectroscopy, and the amount of Mn2+ adsorbed and its maximum adsorption capacity were determined at different pHs. Each set of experiments was repeated three times, and the average value was taken.
| Experimental group | ①1) | ②1) | ③1) | ④1) | ⑤1) | ⑥1) | ⑦2) | ⑧2) | ⑨2) | ⑩2) |
|---|---|---|---|---|---|---|---|---|---|---|
| C(Mn2+) mmol/L | 0 | 0.05 | 0.1 | 0.3 | 0.7 | 1 | 5 | 10 | 25 | 50 |
| V(Mn2+) mL | 0 | 0.5 | 1 | 3 | 7 | 10 | 1 | 2 | 5 | 10 |
| V(DW) mL | 27.5 | 27 | 26.5 | 24.5 | 20.5 | 17.5 | 26.5 | 25.5 | 22.5 | 17.5 |
| Note: 1) The concentration of the added MnSO4 solution was 3 mmol/L; 2) the concentration of the added MnSO4 solution was 0.15 mol/L. | ||||||||||
2.1.4 Kinetics experiments
Briefly, 3 g of wet δ-MnO2 (the weight of the solid sample is 0.624g) was transferred to 200 mL of deionized water containing 25 mM MnSO4. Then the pH of suspension was maintained at ~ 2.5 or 4 for 15 d under vigorous stirring by addition of 0.1 M NaOH or 0.1 M H2SO4. Aliquots were taken at various time intervals and the solids were collected by centrifugation at 16 000g (Neofuge 23R). The collected solids were then rinsed with DDW until the supernatant conductivity was below 20 µS/cm.
2.2 Characterization
2.2.1 X-ray diffraction (XRD) analyses of minerals
The synthesized manganese oxide minerals or intermediates were completely washed with DDW, and then X-ray diffraction (XRD) analyses were carried out with a D8 ADVANCE instrument from the BRUKER Company in Germany. Powder tableting and directional tableting processes were used. The test conditions were as follows: Cu Ka radiation (λ = 0.15406 nm), a 40 kV test voltage, a 40 mA test current, steps, 10°/min scanning speed, and 0.01° step sizes.
2.2.2 Transmission electron microscopy (TEM)
The morphologies and crystalline characteristics of the minerals were analyzed with an FEI-Talos-F200 analytical transmission electron microscope (TEM) and identified by electron diffraction (ED). The acceleration voltage was 200 kV, and the emission current was 10 µA. Ten mL of the sample was added to anhydrous ethanol in a 20 mL capacity serum bottle, small amounts of minerals were added to anhydrous ethanol, and after ultrasonic dispersion, the solution was slightly discolored. Samples were taken from the suspension with an ultrathin carbon membrane for computer analyses.
2.2.3 Flame Atomic Absorption Spectrometry
The concentrations of Mn2+ in the solutions were determined with a Varian AAS240 FS atomic absorption spectrometer. The corresponding elemental lamps and standard curves were used for analyses.
3.1 Mn2+ adsorption at different pH
As shown in Fig. 1a, the amount of Mn2+ adsorbed under acidic conditions increased with the addition of the same amount of Mn2+ and with an increase in pH. When the initial concentration of Mn2+ in the solution was less than 5 mmol/L, δ-MnO2 was completely adsorbed at both pH 2.5 and 4. However, as the concentration of Mn2+ in the solution continued to increase, the adsorption of Mn2+ by δ-MnO2 reached saturation at pH 2.5. At pH 4, the adsorption of Mn2+ by δ-MnO2 reached saturation when the concentration of Mn2+ reached 22.5 mmol/L. This is attributed to the point of zero charge (PZC) of δ-MnO2, which ranges between 1.5 and 2.5. With the increased pH, the negative charge on δ-MnO2 also increased, subsequently enhancing its Mn2+ adsorption capacity. Notably, the maximum adsorption capacity for Mn2+ is lower at pH 2.5 than that at pH 4.
We conducted kinetic experiments using 25 mM Mn2+ with 3 g/L δ-MnO2, and the results are shown in Fig. 1b. When the solution had a pH of 2.5 and 4, the concentration of Mn2+ in the solution rapidly decreased from 25 mM to 2.5 mM and 3.6 mM after the addition, respectively. The kinetic process of δ-MnO2 adsorbing Mn2+ can be divided into two parts. Initially, in the first 48 h, the adsorption of Mn2+ by δ-MnO2 showed minor fluctuations, indicating instability in the adsorption of Mn2+ during this phase. This is likely due to Mn2+ continuously undergoing electron transfer with δ-MnO2 and exchanging with Mn (IV) in bulk, which promotes the crystallization and structural transformation of δ-MnO2.29, 30 Subsequently, in the second part, from 48 hours to 15 days, the amount of Mn2+ adsorbed by δ-MnO2 gradually increased over time. The distinctly different two-stage adsorption process is likely due to the structural and morphological evolution of δ-MnO2 caused by Mn2+.
3.2 XRD patterns of the isothermal adsorption sample
The XRD pattern of the sample obtained after 15 d of reaction is shown in Fig. 2. The XRD pattern of the samples when the Mn2+ concentrations were 0, 1 mmol/L and 5 mmol/L shows two broad diffraction peaks at 37° (d(100) = 0.24 nm) and 65° (d(110) = 0.14 nm) (Fig. 2), which can be attributed to δ-MnO2 with poor crystallinity, small-sized and randomly stacked [MnO6] octahedral layers.21,22 The d-spacing ratio of d(100) to d(110) is 1.73, which indicates a hexagonal layer symmetry.21 When Mn2+ was added up to 10 mmol/L, the peak at d(100) = 0.24 nm began to sharpen and a new shoulder peak appeared at 42.2° 2θ, indicating that Mn2+ was gradually adsorbed on the top and bottom of the octahedral vacancies (Fig. 2). When Mn2+ was added to more than 25 mmol/L, (110), (200), and (310) α-MnO2 peaks appeared (ICDD No. 00-29-1020, d (110) = 0.69 nm, d (200) = 0.48 nm, and d(310) = 0.31 nm) (Fig. 2). In the isothermal adsorption curve, when the concentration of Mn2+ in the solution reached 10 mmol/L, the adsorption capacity of δ-MnO2 for Mn2+ was close to saturation (Fig. 1a). In the XRD patterns, even in this state of adsorption saturation, δ-MnO2 did not undergo structural transformation. As more Mn2+ was added, the amount of Mn2+ adsorbed by δ-MnO2 did not change significantly, but δ-MnO2 transformed into α-MnO2 structure. It might be because the appropriate amount of Mn2+ did not directly integrate into δ-MnO2, but rather facilitated the transformation of δ-MnO2 to α-MnO2 through its effect on the electron transfer of surface Mn (III/IV).
3.3 XRD patterns of the adsorption kinetics samples
The XRD patterns of the obtained samples under pH 2.5 are shown in Fig. 3. The initial mineral was poorly crystalline δ-MnO2. After 2 days of reaction, the intensity of the diffraction peak at d = 0.24 nm increased. And at the same time, the characteristic diffraction peak of α-MnO2 (d (110) = 0.70 nm) appeared at 12° 2θ, which indicated that δ-MnO2 was transformed into α-MnO2 after 2 days (Fig. 2a).7, 11 When the reaction time was extended to 7 days, the intensity of the diffraction peak gradually increased indicated that the δ-MnO2 transformed into α-MnO2 (Fig. 2a). Before the transformation, δ-MnO2 continuously exchanged with Mn2+ in the solution, preparing for the conversion to α-MnO2, such as the generation and arrangement of Mn (III).7 Therefore, the adsorption amount of Mn2+ by δ-MnO2 was constantly fluctuating. When the reaction proceeded for two days, δ-MnO2 transformed into the more stable α-MnO2, and α-MnO2 further adsorbed Mn2+ from the solution.
The XRD patterns of the samples, capturing the dynamic interaction between δ-MnO2 and 25 mM Mn2+ in a pH 4 solution over a period of 15 days, are shown in the Fig. 3b. Compared to the rate at which δ-MnO2 begins to convert to α-MnO2 within 2 days in a pH 2.5 solution, δ-MnO2 in a pH 4 solution only starts to transform into α- MnO2 after 7 d. When the reaction time was 4 days, the peak intensity of d(100) = 0.24 nm increased, indicating that the crystallinity of the δ-MnO2 increased with increasing reaction time.18 When the reaction time was extended to 7 days, the peak intensity at d = 0.24 nm increased, and the characteristic diffraction peak for α-MnO2 (d(110) = 0.70 nm) appeared at 2θ = 12°. This indicated that after 7 days, δ-MnO2 transformed into α-MnO2.7, 11 Therefore, compared to the reaction system at pH 2.5, δ-MnO2 in pH 4 exhibits a fluctuation cycle (preparation period) of 7 days in Mn2+ adsorption. An increase in adsorption is observed when δ-MnO2 begins its transformation into α-MnO2 on the 7 days.
3.4 Morphological changes of δ-MnO2 under different pH
3.4.1 Morphological of Initial δ-MnO2
Figure 4a illustrates the initial formation of nanoflakes aggregates upon mixing the KMnO4 and MnSO4 solutions. Further captured at a higher resolution (Fig. 4b), reveals poorly crystalline δ-MnO2 nanosheets ranging in size from 3 to 5 nm, featuring a lattice spacing of d(100) = 0.24 nm. Furthermore, the selected area diffraction (SAED) pattern (inset of Figure. 4a) shows two diffuse diffraction rings at ∼0.24 nm and ∼0.14 nm, which is consistent with the d(100) = 0.24 nm and d(110) = 0.14 nm spacings of δ-MnO2, respectively. This observation consists with the results obtained from XRD pattern in Fig. 3
3.4.2 Morphological changes of δ-MnO2 under pH 2.5
When δ-MnO2 reacted with Mn2+ for 1 hours at pH 2.5, as demonstrated in Fig. 5a and b, nanoribbons with lengths of 10–20 nm were formed. Upon magnifying nanoribbons, it becomes revealed that these nanoribbons are assembled from several δ-MnO2 nanoflakes. They display lattice with a spacing of d(100) = 0.24 nm, while lacking the typical lattice spacing of d = 0.7 nm characteristic of α-MnO2. This indicates that these nanoribbons are assembled δ-MnO2.
When the reaction time was extended to 1 day, as shown in Fig. 5c and d, the δ-MnO2 nanoflakes gradually disappeared, while the nanoribbons increased in length to 100 nm. The δ-MnO2 nanoflakes were gradually transforming into δ-MnO2 nanoribbons as there were no lattice stripes of α-MnO2 observed. The side perspective of a nanoribbon highlight by yellow dashed line in Fig. 5d demonstrates that an δ-MnO2 nanoflake assemble at the end surface of a nanoribbon.
When the reaction time was extended to 4 days, the δ-MnO2 nanoflakes almost disappeared, resulting in the formation of nanorods with good crystallinity and lengths ranging from 100 nm to 300 nm (Fig. 6a). Upon further magnification of the nanorods, the lattice with a spacing of d(110) = 0.70 nm characteristic of α-MnO2 structures were observed internally (Fig. 6b), consistent with the XRD pattern (Fig. 3a). Small nanorods were assembled side-to-side with each other, increasing the thickness of the original nanorods observed in Fig. 6b. Previous studies also described this side-to-side assembly of α-MnO2 nanorod along the (110) plane driven by surface energy.11, 21
When the reaction time reached 7 days, the nanorods grew to approximately 400 nm and tended to grow in the same direction shown in Fig. 6c. As shown by the amplified image in Fig. 6d, short nanorods with length of about 150 nm connect with each other by end-to-end assembly along the α-MnO2 (001) planes to form secondary nanorods.
3.4.3 Morphological changes of δ-MnO2 under pH 4
When all other conditions remained unchanged and only the pH value was adjusted from 2.5 to 4, nanoribbons with weak crystallinity appeared after a reaction time of 12 h shown in Fig. 7a. The internal lattice of the nanoribbon is inconsistently oriented, with spacing of d(100) = 0.24 nm, indicating that it is a δ-MnO2 structure. These nanoribbons are approximately 40 nm in length, which is longer than the nanoribbons (length about 20 nm) formed after 1 hours at a pH of 2.5. It is noteworthy that the d(100) = 0.24 nm lattice directions within the nanoribbons are inconsistent, which may be due to the δ-MnO2 nanoribbons being newly assembled by δ-MnO2 nanoflakes and not having the opportunity to adjust orientations (Fig. 7b).
When the reaction time reached 4 days, δ-MnO2 nanoflakes gradually transform into nanoribbons with length from 100 to 200 nm (Figs. 7c and d). δ-MnO2 nanoflakes were observed to assemble on the end faces of the nanoribbons (Fig. 7d). This may be due to the outside nanoflakes of the initial δ-MnO2 nanoflakes aggregation were first assembled into long nanoribbons. Then the inner δ-MnO2 nanoflakes were gradually involved as the reaction progressed. Therefore, δ-MnO2 was still present after 4 days. Figure 7d shows several nanosheets stacking and thickening each other, resulting in the formation of nanorods.
After 7 days of reaction, TEM observations revealed that the δ-MnO2 nanoflakes had disappeared, and nanorods approximately 500 nm in length had emerged (Fig. 8a). The SAED with d(110) = 0.70 nm characteristic of α-MnO2 was observed, indicating the beginning of the transformation from δ-MnO2 to α-MnO2 (Fig. 8b). In Fig. 8b, numerous nanoribbons (30 nm in length and 5 nm in width) were observed assembled at the end of nanorods. Notably, these nanoribbons did not exhibit lattice with a spacing of d(110) = 0.70 nm, suggesting they are still of the δ-MnO2 structure. It is probable that these nanoribbons stacked upon each other during the assembly process, ultimately extending along the assembled faces of the nanorods and forming α-MnO2 nanorods.
After 15 days of reaction, there are nanorods with 300–500 nm in length and 30 nm in width were formed (Fig. 8c). Upon further magnification of the nanorods, the lattice with a spacing of d(110) = 0.70 nm characteristic of α-MnO2 structures were observed (Fig. 8d), which was consistent with the XRD findings after 15 d (Fig. 3b). Additionally, it was observed that the nanorods underwent side-to-side assembly along the (110) crystal planes, leading to an increase in width.
2KMnO4 + 3MnSO4 + 2H2O = 5MnO2 + 2H2SO4 + K2SO4 (1)
The morphological evolution process of δ-MnO2 to α-MnO2 in the presence of Mn2+ in solutions with pH 2.5 and 4 is shown in Schematic 1. As shown by the chemical reaction in equation (1), the initial δ-MnO2 was produced by reducing KMnO4 with MnSO4. The δ-MnO2 formed small nanoflakes about 3 to 5 nm in size with hexagonal symmetry. In pH 2.5 or pH 4 solution, δ-MnO2 reacts with Mn2+ concentrations greater than 10 mM to transform into α-MnO2 nanorods. Specifically, δ-MnO2 nanoparticles with sizes ranging from 3 to 5 nm first assembled edge-to-edge along the (100) plane to form longer δ-MnO2 nanoribbons.28, 31 This is consistent with the findings of Liang et al who reported that the morphology evolution preceded the formation of the mineral phase.7 In pH 2.5 solution, secondary δ-MnO2 nanoribbons were assembled along the (110) surface and thickened to form primary α-MnO2 nanorods. Thirdly, primary α-MnO2 nanorods assemble through end-to-end along the (001) plane to form longer nanorods. Adjacent nanorods align side-to-side along the (110) surface, driven by the high surface energy, to form wider nanorods.
Previous studies have shown that the cation adsorption capacity of δ-MnO2 increases with the rise in pH.32, 33 This study also reveals that the amount of Mn2+ adsorbed by δ-MnO2 at pH 4 is 0.5 mmol/g greater than that in a pH 2.5 solution. The slight increase in δ-MnO2 adsorption capacity for Mn2+ is the result of the combined changes in the number of vacancies and the amount of negative charge as the pH increases. A previous study revealed that after aging δ-MnO2 for several years at various pHs, the δ-MnO2 vacancy content was closely related to the pH, with a lower pH leading to the formation of more δ-MnO2 vacancies.17 The vacancies in δ-MnO2, which have a negative charge due to charge deficiency, can strongly adsorb positively charged ions. Therefore, the number of vacancies in δ-MnO2 at pH 2.5 is greater than in a pH 4 solution. Conversely, the PZC of δ-MnO2 is 1.5-2.5, and as the pH of the solution increases from 2.5 to 4, the increase in negative charge leads to stronger adsorption of Mn2+.33 However, the effect of the vacancies was not as significant as that of the PZC, and the amount of Mn2+ adsorbed by δ-MnO2 at pH 4 was greater than in a pH 2.5.
In this study, when there is a small difference in the amount of Mn2+ adsorbed by δ-MnO2, pH is a key factor in controlling the rate of its transformation to the 2×2 tunnel structure of α-MnO2.32, 33 This may be because at pH 2.5, more vacancies are formed directly, and Mn2+ is directly adsorbed onto these vacancies after its addition.17 According to the reaction in equation 2, a large amount of [Mn (III)O6] tunnel walls are formed. These [Mn (III)O6] octahedra include three unsaturated oxygen molecules that combine with H+ to form -OH. When the amount of -OH is large enough, a network of hydrogen bonds (H-bonding) forms between the [Mn (III)O6] octahedra of adjacent nanoribbons and nanoflakes.7 However, in a pH 4 solution, fewer vacancies are formed in δ-MnO2.34 The initial adsorption of Mn2+ and its reaction with δ-MnO2 generate Mn(III), which migrates onto the vacancies due to the Jahn-Teller effect.7 This process is slower than direct adsorption onto vacancies.35 Therefore, at a solution pH of 2.5, δ-MnO2 begins to transform into α-MnO2 within 2 d, while in a pH 4 solution, this process in δ-MnO2 only starts after 7 days.
In comparison to the solution at pH 2.5, the nanoflakes in the pH 4 solution assembled to form longer nanoribbons due to the prolonged transformation process. This is because the formation of [Mn(III)O6] tunnel walls takes longer at pH 4 compared to pH 2.5, allowing δ-MnO2 nanoflakes enough time to assemble edge-to-edge and form longer nanoribbons.7, 36, 37 Therefore, when these long nanoribbons have assembled enough tunnel walls, they directly form primary α-MnO2 nanorods longer than those formed at pH 2.5. When nanoribbons are assembled to form nanorods in a pH 4 solution, many δ-MnO2 nanoflakes remain unassembled into nanoribbons and directly assemble on the (100) crystal face of the primary nanorods. As the pH decreases, the H-bonding between -OH groups on the edges of δ-MnO2 nanoflakes becomes stronger, resulting in the rapid assembly of smaller nanoribbons at pH 2.5. In contrast, when the primary nanorods begin to form at pH 4, some δ-MnO2 nanosheets are still present because the H-bonding is weakened. This phenomenon continues, and subsequently, the uncompleted nanoribbons continue to assemble directly along the (100) crystal face of the nanorods.
Finally, these primary α-MnO2 nanorods assemble end-to-end and side-to-side, increasing in width and length. Yuan et al. calculated the energy of the α-MnO2 (110) facets to be 0.74 J/m2 and the energy of the (001) facet to be 1.17 J/m2. Driven by the increased surface energy, the nanorods connected head-to-tail along the (001) faces and side-to-side via H-bonding at the (110) face, resulting in further thickening and growth of the α-MnO2 nanorods.11
In summary, these results showed that different pH values and different Mn2+ concentrations had significant impacts on the structural transformation and morphological evolution of δ-MnO2 to α-MnO2. Although this study is not exhaustive, it highlights an important area for further researches on the morphological and structural interactions under different conditions. Future studies could involve quantifying the number of vacancies at different pH levels and exploring the interactions of Mn2+ within the system using isotope studies. Additionally, in-situ morphological characterization techniques can be employed to determine the kinetics and mechanisms of the transformation from δ-MnO2 nanoflakes to α-MnO2 nanorods under different pH conditions.
δ-MnO2 is a kind of 3–5 nm nanoparticles consisting of [Mn(III)O6] octahedra and single-layer flaky manganese oxide with hexagonal symmetry and similar biological functions in a supergene environment. The amount of Mn2+ adsorbed by δ-MnO2 increased with increasing pH, indicating that the negative charge on the mineral surface was greater at pH 4 than at pH 2.5, thus the rate of Mn2+ adsorption was greater at pH 4 than at pH 2.5. Different pH values and Mn2+ concentrations strongly influenced the transformation of δ-MnO2 into α-MnO2. As Mn2+ increases and pH decreases from pH 4 to 2.5, the transformation of δ-MnO2 to α-MnO2 nanorods accelerates. With increasing reaction time, δ-MnO2 nanoparticles grew and thickened at the tops and edges of the nanorods through oriented attachment.
Ethics approval and consent to participate
Not applicable.
Competing interests
The authors declare that they have no competing interests.
Funding
This Project was funded by the Youth Project of Yunnan Provincial Basic Research Program under award number 202201AU070178 and Yunnan Agricultural University Research Initiation Project under award number No. F2022-07.
Author Contribution
Liyan Bi and Xinran Liang wrote the main manuscript, Haoran Hu response the questions and modified the manucript in the revision. Lei Wang and Zuran Li did the experiment, Fangdong Zhan, Yonngmei He and Yanqun Zu prepared figure 1-8 and Yuan Li revised the manuscript.
Acknowledgement
The authors thank the Youth Project of Yunnan Provincial Basic Research Program (No. 202201AU070178) and Yunnan Agricultural University Research Initiation Project (No. F2022-07) for the financial support of this research. Liyan Bi and Haoran Hu contributed equally in this study.
Availability of data and materials
The datasets supporting the conclusions of this article are included within the article.
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Schematic 1 is available in the Supplementary Files section.
No competing interests reported.
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Schematic 1. Diagram of crystal growth during the transformation of δ-MnO2 nanosheets to α-MnO2 nanorods at different pH
