δ-MnO2–Mn3O4 Nanocomposite for Photochemical Water Oxidation

Sep 29, 2016 - We consider that the enhanced water oxidation activity could be explained by the active interface between two components. At the phase ...
1 downloads 9 Views 4MB Size
Download PDF
Research Article www.acsami.org

δ‑MnO2−Mn3O4 Nanocomposite for Photochemical Water Oxidation: Active Structure Stabilized in the Interface Zhibin Geng,† Yanxiang Wang,† Jinghai Liu,†,‡ Guangshe Li,† Liping Li,† Keke Huang,† Long Yuan,† and Shouhua Feng*,† †

State Key Laboratory of Inorganic Synthesis and Preparative Chemistry, College of Chemistry, Jilin University, Changchun 130012, People’s Republic of China ‡ Inner Mongolia Key Lab of Chemistry of Natural Products and Synthesis of Functional Molecules, College of Chemistry and Chemical Engineering, Inner Mongolia University for the Nationalities (IMUN), Tongliao 028000, People’s Republic of China S Supporting Information *

ABSTRACT: Pure phase manganese oxides have been widely studied as water oxidation catalysts, but further improvement of their activities is much challenging. Herein, we report an effective method to improve the water oxidation activity by fabricating a nanocomposite of Mn3O4 and δ-MnO2 with an active interface. The nanocomposite was achieved by a partial reduction approach which induced an in situ growth of Mn3O4 nanoparticles from the surface of δ-MnO2 nanosheets. The optimum composition was determined to be 38% Mn3O4 and 62% δ-MnO2 as confirmed by X-ray photoelectron spectra (XPS) and X-ray absorption spectra (XAS). The δ-MnO2−Mn3O4 nanocomposite is a highly active water oxidation catalyst with a turnover frequency (TOF) of 0.93 s−1, which is much higher than the individual components of δ-MnO2 and Mn3O4. We consider that the enhanced water oxidation activity could be explained by the active interface between two components. At the phase interface, weak Mn−O bonds are introduced by lattice disorder in the transition of hausmannite phase to birnessite phase, which provides active sites for water oxidation catalysis. Our study illustrates a new view to improve water oxidation activity of manganese oxides. KEYWORDS: water oxidation, nanocomposite, manganese oxides, active interface, weak Mn−O bond

■

oxide (AMO) with flexible structure is another typical water oxidation catalyst,9,19 while the water oxidation activity of AMO reported hitherto could not catch up with that of nano-Mn2O3.8 The water oxidation activity of manganese oxides is difficult to improve because a weaker Mn−O bond means a more unstable stucture. Therefore, to design an efficient new water oxidation catalyst of manganese oxide, one has to stabilize the weak Mn− O bond without loss of structural stabilization of the target manganese oxide. Nanocomposite consists of different materials usually exhibiting special interface structures between component materials, which results in unique properties and wide potential applications in many fields.20,21 For example, Degussa (Evonik) P25 TiO2, a typical composite of anatase and rutile particles, displays an enhanced photocatalytic performance over pure anatase and rutile due to the interface interaction.22−24 Molecular dynamics simulations have revealed a lattice disorder at anatase−rutile interfaces,25,26 and the rutile TiO6 octahedron structure could even be formed at the anatase side during the phase transition of anatase to rutile. The disordered interfacial

INTRODUCTION Water splitting is recognized as a promising approach to produce low-cost and sustainable hydrogen fuel.1,2 However, the oxidation of water to oxygen is much more difficult than the hydrogen evolution reaction, which limits the application of water splitting.3,4 In nature, water oxidation takes place in the presence of Mn4CaO5 clusters in photosystem II (PS II);5,6 this precedent inspired the study of manganese oxides as water oxidation catalyst.7,8 Many kinds of manganese oxides were discovered to be efficient water oxidation catalysts;9−13 even so, further improving the catalytic activity of these oxides is still highly challenging due to the structure limitation of existing manganese oxides. It is anticipated that the manganese−oxygen bond is the decisive structural feature in determining the water oxidation activity of manganese oxides.14,15 Usually, manganese oxides with a more flexible or weaker Mn−O bond exhibit better activity for water oxidation catalysis.16,17 Owing to the weak Mn−O bond, bixbyite Mn2O3 with Jahn−Teller distorted edgesharing Mn(III)O6 octahedron is indicated to be the best water oxidation catalyst among single manganese oxides.8,18 However, recent improvements on Mn2O3 activity completely originated from the increase of surface area; very few are associated with the structure innovation. In addition, amorphous manganese © 2016 American Chemical Society

Received: August 10, 2016 Accepted: September 29, 2016 Published: September 29, 2016 27825

DOI: 10.1021/acsami.6b09984 ACS Appl. Mater. Interfaces 2016, 8, 27825−27831

Research Article

ACS Applied Materials & Interfaces Table 1. Mn Oxidation States, BET Surface Areas, and TOFs of the Samples sample

reductant conc

Mn oxidation state

SBETb

molecular formulac

TOFd

0.25 M 0.75 M 1.25 M 2.5 M

4.1 3.7 3.5 3.3 2.7

96 110 112 97 44

K0.05MnO2·2H2O MnO1.9·2.7H2O MnO1.8·2.7H2O MnO1.6·2.9H2O Mn3O4·6.3H2O

0.02 0.53 0.93 0.76 0.39

δ-MnO2a MnOx-1 MnOx-2 MnOx-3 nano-Mn3O4 a

δ-MnO2 was an annealed sample. bSurface area determined by BET measurement (m2 g−1). cMolecular formula calculated from ICP-AES data. Turnover frequency (TOF) of oxygen molecules rate per catalytic sites (mmolO2 molMn−1 s−1).

d

A 0.2 g portion of the as-prepared δ-MnO2 was ultrasonicated to disperse in 50 mL of deionized water. Then, 20 mL portions of NaBH4 aqueous solution with concentrations of 0.25, 0.75, and 1.25 M were added to the suspensions, respectively. The suspensions were stirred for 30 min at room temperature. The obtained samples were washed with deionized water for several times, dried in a lyophilizer, and annealed at 200 °C for 5 h. The three products are named MnOx-1, MnOx-2, and MnOx-3 (Table 1). In order to obtain a nano-Mn3O4 sample with the same morphology as a Mn3O4 nanoparticle in a nanocomposite, an overdose reductant (2.5 M NaBH4) was added to reduce 0.2 g of as-prepared δ-MnO2; other operation procedures followed that mentioned above. A portion of δ-MnO2 nanosheets was also heated at 200 °C for 5 h for comparison. Photochemical Water Oxidation Tests. A typical Ru(bpy)32+− Na2S2O8 system was used to characterize photochemical water oxidation properties.30 A 5 mg portion of the as-prepared sample was added to a photolysis vessel containing 20 mL of Na2SiF6− NaHCO3 buffer with 2 mM Ru(bpy)32+ and 10 mM Na2S2O8; the concentration of sample was 250 ppm. The vessel was illuminated with a 300 W xenon lamp with filters for UV (420 nm cutoff). Dissolved oxygen concentration was measured by a Clark-type oxygen electrode at 20 °C controlled by water recirculator.

region stabilizes a higher charge density and further enhances the photocatalytic activity.26 Thus, designing and fabricating a composite with an active interface of specific structure and electronic states can improve the catalytic activity of materials. Although the catalytic mechanisms of photocatalysis and water oxidation are different, the structure advantage of anatase−rutile TiO2 is worthy of imitation. Manganese oxides have many structural similarities to titanic oxides. The basic structure units of most manganese oxides are MnO 6 octahedra,16 just as the TiO6 octahedra in anatase and rutile. More importantly, the disordered structure widely exists in manganese oxides due to the rich linkage modes of MnO6 octahedra and the variety of Mn oxidation states.27,28 Thus, manganese oxide nanocomposite is predicted to have a special disordered interface, and the weak Mn−O bonds at the disordered interface could improve the water oxidation activity substantially. Moreover, we proposed that the weak Mn−O bonds at the interface might be stabilized in the transition of two different lattices; then, a stable manganese oxide with a much more unstable Mn−O bond can be obtained. Hausmannite Mn3O4 and birnessite δ-MnO2 are both manganese oxides with edge-sharing MnO6 octahedra, and their water oxidation activities are much lower than that of Mn2O3. However, when Mn3O4 and δ-MnO2 are combined together to form a composite, it is highly possible to form much weak Mn−O bonds just like that of anatase−rutile TiO2, which may give rise to an enhanced water oxidation activity. In this work, we utilized a facile reduction approach that induced hausmannite Mn3O4 nanoparticles with in situ growth from pure birnessite δ-MnO2 nanosheets to form the δ-MnO2− Mn3O4 nanocomposite. The nanocomposite showed a much better water oxidation activity than pure δ-MnO2 and Mn3O4. The enhanced activity should be related to the weak Mn−O bond at the phase interface that provides active sites for water oxidation.

■

■

RESULTS AND DISCUSSION XRD was used to identify the crystallinity and structure of the as-synthesized manganese oxide samples, and corresponding SEM images were given on the right (Figure 1). XRD patterns of δ-MnO2 nanosheets compare well with standard data of the birnessite phase.9,31 Crystal facets (001), (002), (100), and

EXPERIMENTAL SECTION

Materials. Chemicals KMnO4 (99.5%), NaHCO3 (99.5%), Na2SiF6 (99%), Na2S2O8 (98%), ethyl acetate (99.5%), and NaBH4 (96%) were purchased from Sinopharm Chemical Reagent Co. Ltd. [Ru(bpy)3]Cl2·6H2O (98%) was purchased from J&K Scientific Ltd. Commercial Mn2O3 (99%, ∼325 mesh) was purchased from SigmaAldrich, and Mn3O4 (M111159, Mn ≥ 71.0%, BET surface area 5−7 m2/g) was from Aladdin. All reagents and solvents were used as received. Synthesis Procedures. δ-MnO2 nanosheets were prepared, just following the procedure reported by Pal at al.29 Namely, 3.16 g of KMnO4 was dissolved in 750 mL of deionized water, and the solution was moved into a 1 L round-bottom flask. Then, 200 mL of ethyl acetate was added into the above KMnO4 solution to form a biphasic system. The flask was kept in a water bath with refluxing devices at 80−85 °C for 10 h; brown δ-MnO2 precipitates were formed at the bottom of the flask. The products were washed with deionized water several times, and dried in a lyophilizer.

Figure 1. Powder XRD patterns and SEM images of δ-MnO2, MnOx1, MnOx-2, and MnOx-3. Standard XRD data for hausmannite Mn3O4 (PDF 24-0734) was put beneath for comparison. Yellow arrows in SEM image of MnOx-1 pointed to nanoparticles. 27826

DOI: 10.1021/acsami.6b09984 ACS Appl. Mater. Interfaces 2016, 8, 27825−27831

Research Article

ACS Applied Materials & Interfaces (110) appeared at d-spacings of 7.04 Å (11.95°), 3.68 Å (24.18°), 2.42 Å (37.06°), and 1.41 Å (66.08°), respectively. Weak peak intensities were indicative of lower crystallinity of ultrathin 2-dimensional materials than bulk materials.29 SEM images showed that δ-MnO2 was folded as a 2-dimensional nanosheet, in keeping with the results of Pal at al.29 TEM images of δ-MnO2 were shown in Figure S1. XRD patterns of MnOx-1, MnOx-2, and MnOx-3 showed an increase in the Mn3O4 (hausmannite, PDF 24-0734) phase, as followed by a decrease in δ-MnO2 phase (Figure 1). MnOx-1 is mainly composed of δ-MnO2 phase, while peaks of (001) and (002) planes faded, suggesting a partial collapse of the layered structure. A broad peak appeared at around 2θ of 53°, indicating that some interlayer Mn(III)O6 octahedron structures were formed.32 MnOx-2 was a mixed-phase material that contained Mn3O4 and δ-MnO2 phases, while MnOx-3 mainly consisted of the Mn3O4 phase. This result indicates that δ-MnO2 tends to be reduced to Mn3O4 with the dosage augment of reductant. IR graphs further confirmed this conclusion (Figure S2). For the SEM images (full images in Figure S4), nanoparticles with diameter of 30−40 nm gradually appeared, while the nanosheets gradually disappeared from MnOx-1 to MnOx-3. Few nanoparticles were found in MnOx-1, and MnOx-2 consisted of a merely equal ratio of nanosheets and nanoparticles, while MnOx-3 was constructed by mostly nanoparticles. The progress of nanoparticles increasing and nanosheets wearing off corresponds to the changing trend of the δ-MnO2 and Mn3O4 phases shown in XRD patterns (Figure 1). XRD and SEM data for nano-Mn3O4 were shown in Figure S6. The morphology of nano-Mn3O4 is in accordance with that of a nanoparticle in MnOx-2 and MnOx-3, indicating that the nanoparticle was Mn3O4. The detailed morphology and composition of MnOx-2 were studied by HRTEM. As shown in Figure 2a, the nanoparticles distributed evenly on the surface and inside the folded structure of nanosheets, in accordance with that observed by the SEM

image for MnO x-2. The nanosheets and nanoparticles combined so firmly with each other that an individual nanoparticle or nanosheet could not be found. With a zoom in on the red square of Figure 2a, Mn3O4 nanoparticles are clearly shown (Figure 2b). Since single nanoparticles did not exist, fast Fourier transform (FFT) was used to define the structure of the crystal. FFT results indicate the presence of (321) and (211) planes of Mn3O4. The interface of the nanosheet and nanoparticle is shown in Figure 2c, which is the orange square in Figure 2b. The nanosheets are concluded to be δ-MnO2 as the d-spacing of around 2.43 Å for (100) is shown. Among the interface, we incidentally found that a piece of the δ-MnO2 nanosheet is partly covered over the Mn3O4 nanoparticle, and their lattice fringes finally overlapped together (green square in Figure 2c). This phenomenon may reflect a formation mechanism of the nanocomposite. That is, during the reduction progress, some interlayer MnO6 octahedra gradually occurred, and then layers of δ-MnO2 nanosheets at this position collapsed and formed Mn3O4 nuclei; eventually, the nuclei grew and expanded to Mn3O4 nanoparticles. As a consequence, some δ-MnO2 nanosheets were covered over Mn3O4 nanoparticles in the expansion progress. Figure 2d shows the detailed lattice of the interface, which is the green square of Figure 2c. One can see that the lattice fringes of δMnO2 are not parallel at the interface, as demonstrated by the yellow arrows. The last few MnO6 octahedra of δ-MnO2 near the interface are obviously influenced by the Mn3O4 lattice, and thus become distorted. This indicates that the interface of the δ-MnO2−Mn3O4 nanocomposite is lattice disordered. Brunauer−Emmett−Teller (BET) surface areas of δ-MnO2, MnOx-1, MnOx-2, MnOx-3, and nano-Mn3O4 were estimated to be 96, 110, 112, 97, and 44 m2 g−1, respectively (Table 1). Comparing these with the electron microscope images of these samples, we concluded that the high surface area of the nanocomposite mainly originated from δ-MnO2 nanosheets, while nano-Mn3O4 also showed a relatively high BET surface. Oxidation states of Mn ions in δ-MnO2 nanosheets and manganese oxide catalysts were analyzed by XPS. Binding energy of Mn 2p and Mn 3s were shown in Figure 3. Mn 2p1/2 and Mn 2p3/2 peaks of δ-MnO2 nanosheets were at 654.3 and 642.7 eV. The shape and binding energy of Mn 2p3/2 peaks are consistent with those reported for birnessite MnO2.33,34 For the

Figure 2. HRTEM images of MnOx-2: (a) general view; (b) Mn3O4 nanoparticles (red square in panel a), and fast Fourier transform (FFT) of nanoparticle district; (c) interface of δ-MnO2 nanosheets and Mn3O4 nanoparticles (orange square in panel b); and (d) detailed lattice of the interface (green square in panel c). The lattice fringes of δ-MnO2 are not parallel at the interface.

Figure 3. XPS of δ-MnO2, MnOx-1, MnOx-2, and MnOx-3. (a) Mn 2p spectra. (b) Mn 3s spectra. The variations in the XPS data indicate the changes of Mn oxidation states of samples. 27827

DOI: 10.1021/acsami.6b09984 ACS Appl. Mater. Interfaces 2016, 8, 27825−27831

Research Article

ACS Applied Materials & Interfaces Mn 2p3/2 spectrum, a characteristic peak of Mn(III) locates at 641.0 ± 0.2 eV; no peaks are observed in our spectrum at this binding energy, which implies the absence of Mn(III).33 Therefore, the oxidation state of manganese was around +4. In contrast, the binding energies of MnOx-1, MnOx-2, and MnOx3 were lower than those of δ-MnO2 nanosheets. An obvious peak appeared at 641 ± 0.2 eV, and its area ratio increased from MnOx-1, to MnOx-2, to MnOx-3, indicating that the oxidation state of Mn ions in MnOx-1, MnOx-2, and MnOx-3 gradually decreased (Figure 3a). However, XPS of Mn 2p is very complicated since the Mn 2p spectrum of every single Mn oxidation state actually consists of several peaks.33,34 Thus, accurate oxidation states cannot be obtained by Mn 2p spectra. Meanwhile, the particular multiplet splitting of Mn 3s peaks was capable of distinguishing Mn oxidation states.35,36 The Mn 3s spectrum has two peaks caused by the coupling of nonionized 3s electron with 3d valence-band electrons; the magnitude of peak splitting is in proportion to the oxidation state. According to the data reported by V. R. Galakhov et al.,37 the energy separation of Mn 3s peaks from Mn(IV) oxides to Mn(III) oxides was 4.4−5.3 eV. For comparison, the energy separations of Mn 2s for δ-MnO2, MnOx-1, MnOx-2, and MnOx-3 were 4.33, 4.64, 4.85, and 5.07 eV (Figure 3b), suggesting the Mn oxidation states of 4.1, 3.7, 3.5, and 3.3 for these samples in sequence (Table 1). From the XPS data, it is shown that the Mn oxidation state decreased gradually along with the increase of reductant dosage, which conforms to the composition change of samples. As the oxidation states of δ-MnO2 and Mn3O4 are settled, the theoretical composition of samples could be calculated from the average oxidation state of manganese. The theoretical ratio of δ-MnO2 and Mn3O4 phase in MnOx-1, MnOx-2 and MnOx-3 was 0.80:0.20, 0.62:0.38, and 0.44:0.56, respectively (calculated in Mn ion ratio). Compared to SEM images, the calculated components ratios of MnOx-2 and MnOx-3 corresponded to their morphologies, but MnOx-1 obviously contains less than 20% Mn3O4. We considered that, in MnOx-1, interlayer Mn(III) caused the decrease of the Mn oxidation state, which could be the crystal seed for Mn3O4 nanoparticles when the reduction intensity was higher. Mn K-edge spectra of δ-MnO2, MnOx-2, and commercial Mn3O4 were comparatively studied to analyze the detailed structure and chemical information (Figure 4). The XANES measurements, shown in Figure 4a, followed the trend of oxidation state changes observed from XPS characterization. When δ-MnO2 was reduced to MnOx-2, an obvious negative shift was observed in the edge position, indicating the reduction of the Mn oxidation state. With reference to the edge position and shape of commercial Mn3O4 XANES, MnOx-2 was a mixed-phase material consisting of δ-MnO2 and Mn3O4. From the linear combination fitting, the composition of MnOx-2 was determined to contain 40.9% Mn3O4 and 59.1% δ-MnO2, significantly closer to the XPS data analysis. It is noteworthy that the fitting curve was perfectly accordant with the actual curve of MnOx-2 with an R-factor of 0.000 174. Therefore, MnOx-2 had no other phases except for Mn3O4 and δ-MnO2. To further confirm the structure and composition of MnOx2, EXAFS data were studied (Figure 4b). Just like the phase composition proved by XRD and XANES, the EXAFS spectrum of MnOx-2 was evidently composed of δ-MnO2 and Mn3O4. The MnOx-2 spectrum showed all characteristic peaks of two phases, and the intensity of every peak was weaker than that of the individual phase. Nevertheless, MnOx-2 showed an

Figure 4. (a) Comparison of XANES collected on δ-MnO2, MnOx-2, and commercial Mn3O4. The blank curve was the linear combination fitting curve of MnOx-2; it was boldfaced for visual purposes. (b) EXAFS of δ-MnO2, MnOx-2, and commercial Mn3O4.

abnormally weak EXAFS peak in the R′ = 3.1 Å region, despite the peak shape that corresponds to the intermediate state of δMnO2 and Mn3O4. This can be explained by smaller particle size of MnOx-2, which was opposed to the extended crystal structure in 2-dimensional δ-MnO2 and commercial Mn3O4.38 Simultaneously, the peak of the R′ = 5 Å region in MnOx-2 EXAFS was much weaker than those for both δ-MnO2 and commercial Mn3O4, which likely originated from the Mn− Mn−Mn multiple scattering,39 suggesting that MnOx-2 has a smaller crystal size in each phase.38 Thus, XAS spectra confirmed the conclusion of XRD and TEM, demonstrating the as-prepared MnOx-2 sample was a mixture phase of δMnO2 and Mn3O4. The only structural difference between the δ-MnO2−Mn3O4 nanocomposite and δ-MnO2−Mn3O4 mechanical mixture lies on the interface structure. Photochemical water oxidation experiments were conducted in buffered deoxygened aqueous solution (pH = 5.8), with Ru(bpy)32+ as photosensitizer and Na2S2O8 as a two-electron acceptor. The catalytic activities of all samples were tested under visible light illumination. The δ-MnO2 was the annealed sample for comparison. As shown in Figure 5, TOFs of all samples were measured at the criterion of oxygen evolution rate per catalysis center (mmolO2 molMn−1 s−1), and the TOF values were calculated from the data of 1−2 min of the oxygen evolution curves. The blank was the control group, indicating that the Ru(bpy)32+−Na2S2O8 system could not produce oxygen without catalyst. The results of photochemical water oxidation measurements are listed in Table 1. MnOx-2 showed the highest activity of 0.93 s−1, better than 0.53 s−1 for MnOx-1 and 0.76 s−1 for MnOx-3, while TOFs of δ-MnO2 (0.02 s−1) and nano-Mn3O4 (0.39 s−1) are much lower. The oxygen 27828

DOI: 10.1021/acsami.6b09984 ACS Appl. Mater. Interfaces 2016, 8, 27825−27831

Research Article

ACS Applied Materials & Interfaces

weak and flexible Mn−O bond is in favor of water oxidation; the unstable MnO6 octahedra at the interface are much more active than MnO6 octahedra in bulk δ-MnO2 and Mn3O4. In addition, both δ-MnO2 and Mn3O4 lattices beside the interface are stable, so that the weak Mn−O bonds at the interface are stabilized in the transition of two different lattices. Thus, a stable manganese oxide with a weak Mn−O bond is formed. Simultaneously, manganese ions at the interface should have a mixed valence of Mn(III) and Mn(IV) due to different Mn oxidation states of δ-MnO2 and Mn3O4. This electronic state of the nanocomposite interface may contribute to high water oxidation activity as the manganese ions in the PSII cluster are also the mixed valence of Mn(III) and Mn(IV).41,42 Thus, the high activity of mixed-phase nanocomposite samples should be caused by a weak Mn−O bond and the mixed valence of Mn at the phase interface of δ-MnO 2 nanosheet and Mn 3 O 4 nanoparticle, which conforms to the structural character of the PSII cluster. A simplified mechanism schematic illustration is shown in Figure 6. Figure 5. Photochemical water oxidation tests of δ-MnO2, MnOx-1, MnOx-2, MnOx-3, and commercial Mn3O4. The black curve was a blank test without a sample.

evolution rates follow such a sequence: MnOx-2 > MnOx-3 > MnOx-1 > nano-Mn3O4 > δ-MnO2. In addition, TOF of MnOx2 was 7.1 times that of commercial Mn2O3 (0.13 s−1), as shown in Figure S5. It is well-documented that δ-MnO2 and Mn3O4 were not high-activity manganese oxides in comparison to Mn2O3.16 However, when δ-MnO2 and Mn3O4 are combined to form a nanocomposite, an activity much higher than individual phase is obtained, which is even comparable to that of nano-Mn2O3.18 It is worth investigating why the nanocomposite is highly active. Obviously, the simple coexistence of two phases is not the reason for high activity. A mechanical mixture of δ-MnO2 and nano-Mn3O4 with molar ratio corresponding to that of MnOx-2 was prepared. The TOF value of the mechanical mixture was merely between δ-MnO2 and nano-Mn3O4, which does not match that for MnOx-2 (Figure S7). The interlayer Mn(III)O6 octahedra could be the alternative active centers for water oxidation, which may widely exist in MnOx-1, but Mn(III)O6 octahedra are located inside the Mn(IV)O6 octahedra layers; this structure prevents the contact of active centers and water, leading to a decrease in overall activity for water oxidation.40 As a result, TOF of MnOx-1 was much lower than the nanocomposite material of MnOx-2. The distinction of mixed-phase nanocomposite samples (MnOx-2, MnOx-3) and mechanical mixture laid on the combination mode; Mn3O4 nanoparticles were grown in situ from δ-MnO2 nanosheets. SEM and TEM images clearly show that δ-MnO2 nanosheets and Mn3O4 nanoparticles are tightly bonded together. We consider that the lattice of δ-MnO2 nanosheets and Mn3O4 nanoparticles directly bond with each other at the interface. The basic structural units of δ-MnO2 and Mn3O4 are both edge-sharing MnO6 octahedron, which is the foundation of the structural combination. However, the structural parameters of their MnO6 octahedra are different: the Mn(III)O6 octahedron in Mn3O4 is Jahn−Teller elongated, while the Mn(IV)O6 octahedron in δ-MnO2 is much more regular. Thus, similar to the anatase−rutile interface,26 a highly disordered structure will be able at the interface to show weak and flexible Mn−O bonds. According to Dismukes et al.,16 a

Figure 6. Schematic illustration of the mixed-phase nanocomposite for photochemical water oxidation. Light gray, gray, and dark gray balls represent Mn(IV), Mn(III), and Mn(II) ions. Red and green balls represent oxygen and hydrogen atoms.

MnOx-3 was a nanocomposite material as well. MnOx-3 contains 56% Mn3O4 phase as compared to that of 38% for MnOx-2. MnOx-3 contained more Mn3O4 phase, while its water oxidation activity was lower than that of MnOx-2. As observed from SEM images, Mn3O4 nanoparticles in MnOx-3 were more separated from δ-MnO2 nanosheets. As the reduction intensity of MnOx-3 was stronger than MnOx-2, we consider that, during the fabrication process of MnOx-3, more Mn3O4 crystal seeds formed and grew into nanoparticles. As a result, many δ-MnO2 nanosheets that were bonded with Mn3O4 nanoparticles transformed into Mn3O4 phase completely, leading to a decrease in interface active centers, and therefore a declined water oxidation activity.

■

CONCLUSIONS New manganese oxide nanocomposites consisting of Mn3O4 nanoparticles and δ-MnO2 nanosheets were fabricated via a facile partial reduction method. Mn3O4 nanoparticles were grown in situ from δ-MnO2 nanosheets when the δ-MnO2 nanosheets precursor was partially reduced. The photochemical water oxidation activity of nanocomposites was far beyond those of the components, Mn3O4 or δ-MnO2. Moreover, the nanocomposite with a composition of 38% Mn3O4 and 62% δMnO2 gave the highest TOF value of 0.93 s−1, which is 27829

DOI: 10.1021/acsami.6b09984 ACS Appl. Mater. Interfaces 2016, 8, 27825−27831

Research Article

ACS Applied Materials & Interfaces

(8) Menezes, P. W.; Indra, A.; Littlewood, P.; Schwarze, M.; Gobel, C.; Schomacker, R.; Driess, M. Nanostructured Manganese Oxides as Highly Active Water Oxidation Catalysts: A Boost from Manganese Precursor Chemistry. ChemSusChem 2014, 7, 2202−2211. (9) Iyer, A.; Del-Pilar, J.; King’ondu, C. K.; Kissel, E.; Garces, H. F.; Huang, H.; El-Sawy, A. M.; Dutta, P. K.; Suib, S. L. Water Oxidation Catalysis using Amorphous Manganese Oxides, Octahedral Molecular Sieves (OMS-2), and Octahedral Layered (OL-1) Manganese Oxide Structures. J. Phys. Chem. C 2012, 116, 6474−6483. (10) Wiechen, M.; Najafpour, M. M.; Allakhverdiev, S. I.; Spiccia, L. Water Oxidation Catalysis by Manganese Oxides: Learning from Evolution. Energy Environ. Sci. 2014, 7, 2203−2212. (11) Pokhrel, R.; Goetz, M. K.; Shaner, S. E.; Wu, X.; Stahl, S. S. The “Best Catalyst” for Water Oxidation Depends on the Oxidation Method Employed: A Case Study of Manganese Oxides. J. Am. Chem. Soc. 2015, 137, 8384−8387. (12) Boppana, V. B.; Jiao, F. Nanostructured MnO2: an Efficient and Robust Water Oxidation Catalyst. Chem. Commun. 2011, 47, 8973− 8975. (13) Smith, P. F.; Deibert, B. J.; Kaushik, S.; Gardner, G.; Hwang, S.; Wang, H.; Al-Sharab, J. F.; Garfunkel, E.; Fabris, L.; Li, J.; Dismukes, G. C. Coordination Geometry and Oxidation State Requirements of Corner-Sharing MnO6 Octahedra for Water Oxidation Catalysis: An Investigation of Manganite (γ-MnOOH). ACS Catal. 2016, 6, 2089− 2099. (14) Jeong, D.; Jin, K.; Jerng, S. E.; Seo, H.; Kim, D.; Nahm, S. H.; Kim, S. H.; Nam, K. T. Mn5O8 Nanoparticles as Efficient Water Oxidation Catalysts at Neutral pH. ACS Catal. 2015, 5, 4624−4628. (15) Shen, J. R. The Structure of Photosystem II and the Mechanism of Water Oxidation in Photosynthesis. Annu. Rev. Plant Biol. 2015, 66, 23−48. (16) Robinson, D. M.; Go, Y. B.; Mui, M.; Gardner, G.; Zhang, Z.; Mastrogiovanni, D.; Garfunkel, E.; Li, J.; Greenblatt, M.; Dismukes, G. C. Photochemical Water Oxidation by Crystalline Polymorphs of Manganese Oxides: Structural Requirements for Catalysis. J. Am. Chem. Soc. 2013, 135, 3494−3501. (17) Indra, A.; Menezes, P. W.; Driess, M. Uncovering Structure− Activity Relationships in Manganese-Oxide-Based Heterogeneous Catalysts for Efficient Water Oxidation. ChemSusChem 2015, 8, 776−785. (18) Kuo, C.-H.; Mosa, I. M.; Poyraz, A. S.; Biswas, S.; El-Sawy, A. M.; Song, W.; Luo, Z.; Chen, S.-Y.; Rusling, J. F.; He, J.; Suib, S. L. Robust Mesoporous Manganese Oxide Catalysts for Water Oxidation. ACS Catal. 2015, 5, 1693−1699. (19) Indra, A.; Menezes, P. W.; Zaharieva, I.; Baktash, E.; Pfrommer, J.; Schwarze, M.; Dau, H.; Driess, M. Active Mixed-Valent MnOx Water Oxidation Catalysts through Partial Oxidation (Corrosion) of Nanostructured MnO Particles. Angew. Chem., Int. Ed. 2013, 52, 13206−13210. (20) Xue, P.; Yang, X.; Lai, X.; Xia, W.; Li, P.; Fang, J. Controlling Synthesis and Gas-Sensing Properties of Ordered Mesoporous In2O3Reduced Graphene Oxide (rGO) Nanocomposite. Sci. Bull. 2015, 60, 1348−1354. (21) Hu, H.; Guan, B.; Xia, B.; Lou, X. W. Designed Formation of Co3O4/NiCo2O4 Double-Shelled Nanocages with Enhanced Pseudocapacitive and Electrocatalytic Properties. J. Am. Chem. Soc. 2015, 137, 5590−5595. (22) Lin, H.; Li, L.; Zhao, M.; Huang, X.; Chen, X.; Li, G.; Yu, R. Synthesis of High-Quality Brookite TiO2 Single-Crystalline Nanosheets with Specific Facets Exposed: Tuning Catalysts from Inert to Highly Reactive. J. Am. Chem. Soc. 2012, 134, 8328−8331. (23) Hirakawa, H.; Katayama, M.; Shiraishi, Y.; Sakamoto, H.; Wang, K.; Ohtani, B.; Ichikawa, S.; Tanaka, S.; Hirai, T. One-Pot Synthesis of Imines from Nitroaromatics and Alcohols by Tandem Photocatalytic and Catalytic Reactions on Degussa (Evonik) P25 Titanium Dioxide. ACS Appl. Mater. Interfaces 2015, 7, 3797−3806. (24) Li, G.; Gray, K. A. The Solid-Solid Interface: Explaining the High and Unique Photocatalytic Reactivity of TiO2-Based Nanocomposite Materials. Chem. Phys. 2007, 339, 173−187.

comparable to nano-Mn2O3. The reason for enhanced catalytic activity observed from the new nanocomposites could be attributed to the weak Mn−O bond at the interface due to the disordered structure in the lattice transition from Mn3O4 to δMnO2. The stable lattice of Mn3O4 and δ-MnO2 besides the interface stabilized the disordered structure and formed a stable material. The results reported in this work provide a new method for improving the water oxidation activity of manganese oxides. Since the structure and exposed area of active interface can be further optimized, nanocomposite material still has huge potential for water oxidation.

■

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acsami.6b09984. Additional experimental results, detailed SEM and TEM images of samples, other comparisons of water oxidation activity, and the influence of annealing treatment for the samples (PDF)

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This work was supported by the National Natural Science Foundation of China (Grants 21427802, 21131002, 21201075, 21671076, and 21303080) and the Specialized Research Fund for the Doctoral Program of Higher Education (SRFDP Grant 20110061130005).

■

REFERENCES

(1) Wu, H. B.; Xia, B. Y.; Yu, L.; Yu, X. Y.; Lou, X. W. Porous Molybdenum Carbide Nano-Octahedrons Synthesized via Confined Carburization in Metal-Organic Frameworks for Efficient Hydrogen Production. Nat. Commun. 2015, 6, 6512−6519. (2) Yu, X. Y.; Yu, L.; Wu, H. B.; Lou, X. W. Formation of Nickel Sulfide Nanoframes from Metal-Organic Frameworks with Enhanced Pseudocapacitive and Electrocatalytic Properties. Angew. Chem., Int. Ed. 2015, 54, 5331−5335. (3) Murakami, M.; Hong, D.; Suenobu, T.; Yamaguchi, S.; Ogura, T.; Fukuzumi, S. Catalytic Mechanism of Water Oxidation with Single-Site Ruthenium-Heteropolytungstate Complexes. J. Am. Chem. Soc. 2011, 133, 11605−11613. (4) Han, L.; Yu, X. Y.; Lou, X. W. Formation of Prussian-Blue-Analog Nanocages via a Direct Etching Method and their Conversion into Ni−Co-Mixed Oxide for Enhanced Oxygen Evolution. Adv. Mater. 2016, 28, 4601−4605. (5) Petrie, S.; Pace, R. J.; Stranger, R. Resolving the Differences Between the 1.9 Å and 1.95 Å Crystal Structures of Photosystem II: A Single Proton Relocation Defines Two Tautomeric Forms of the Water-Oxidizing Complex. Angew. Chem., Int. Ed. 2015, 54, 7120− 7124. (6) Walsh, D.; Sanchez-Ballester, N. M.; Ariga, K.; Tanaka, A.; Weller, M. Chelate Stabilized Metal Oxides for Visible Light Photocatalyzed Water Oxidations. Green Chem. 2015, 17, 982−990. (7) Robinson, D. M.; Go, Y. B.; Greenblatt, M.; Dismukes, G. C. Water Oxidation by λ-MnO2: Catalysis by the Cubical Mn4O4 Subcluster Obtained by Delithiation of Spinel LiMn2O4. J. Am. Chem. Soc. 2010, 132, 11467−11469. 27830

DOI: 10.1021/acsami.6b09984 ACS Appl. Mater. Interfaces 2016, 8, 27825−27831

Research Article

ACS Applied Materials & Interfaces (25) Deskins, N. A.; Kerisit, S.; Rosso, K. M.; Dupuis, M. Molecular Dynamics Characterization of Rutile-Anatase Interfaces. J. Phys. Chem. C 2007, 111, 9290−9298. (26) Nolan, M.; Deskins, N. A.; Schwartzenberg, K. C.; Gray, K. A. Local Interfacial Structure Influences Charge Localization in Titania Composites: Beyond the Band Alignment Paradigm. J. Phys. Chem. C 2016, 120, 1808−1815. (27) Dose, W. M.; Sharma, N.; Webster, N. A. S.; Peterson, V. K.; Donne, S. W. Kinetics of the Thermally-Induced Structural Rearrangement of γ-MnO2. J. Phys. Chem. C 2014, 118, 24257−24265. (28) Chabre, Y.; Pannetier, J. Structural and Electrochemical Properties of the Proton/γ-MnO2 System. Prog. Solid State Chem. 1995, 23, 1−130. (29) Sinha, A. K.; Pradhan, M.; Pal, T. Morphological Evolution of Two-Dimensional MnO2 Nanosheets and Their Shape Transformation to One-Dimensional Ultralong MnO2 Nanowires for Robust Catalytic Activity. J. Phys. Chem. C 2013, 117, 23976−23986. (30) Morris, N. D.; Suzuki, M.; Mallouk, T. E. Kinetics of Electron Transfer and Oxygen Evolution in the Reaction of [Ru(bpy)3]3+ with Colloidal Iridium Oxide. J. Phys. Chem. A 2004, 108, 9115−9119. (31) Polzer, F.; Kunz, D. A.; Breu, J.; Ballauff, M. Formation of Ultrathin Birnessite-Type Nanoparticles Immobilized on Spherical Polyelectrolyte Brushes. Chem. Mater. 2010, 22, 2916−2922. (32) Drits, V. A.; Lanson, B.; Gaillot, A. C. Birnessite Polytype Systematics and Identification by Powder X-ray Diffraction. Am. Mineral. 2007, 92, 771−788. (33) Banerjee, D.; Nesbitt, H. W. XPS Study of Dissolution of Birnessite by Humate with Constraints on Reaction Mechanism. Geochim. Cosmochim. Acta 2001, 65, 1703−1714. (34) Biesinger, M. C.; Payne, B. P.; Grosvenor, A. P.; Lau, L. W. M.; Gerson, A. R.; Smart, R. S. C. Resolving Surface Chemical States in XPS Analysis of First Row Transition Metals, Oxides and Hydroxides: Cr, Mn, Fe, Co and Ni. Appl. Surf. Sci. 2011, 257, 2717−2730. (35) Foord, J. S.; Jackman, R. B.; Allen, G. C. An X-ray Photoelectron Spectroscopic Investigation of the Oxidation of Manganese. Philos. Mag. A 1984, 49, 657−663. (36) Wertheim, G. K.; Hüfner, S.; Guggenheim, H. J. Systematics of Core-Electron Exchange Splitting in 3d-Group Transition-Metal Compounds. Phys. Rev. B 1973, 7, 556−558. (37) Galakhov, V. R.; Demeter, M.; Bartkowski, S.; Neumann, M.; Ovechkina, N. A.; Kurmaev, E. Z.; Lobachevskaya, N. I.; Mukovskii, Y. M.; Mitchell, J.; Ederer, D. L. Mn 3s exchange splitting in mixedvalence manganites. Phys. Rev. B: Condens. Matter Mater. Phys. 2002, 65, 113102. (38) Gorlin, Y.; Lassalle-Kaiser, B.; Benck, J. D.; Gul, S.; Webb, S. M.; Yachandra, V. K.; Yano, J.; Jaramillo, T. F. In Situ X-ray Absorption Spectroscopy Investigation of a Bifunctional Manganese Oxide Catalyst with High Activity for Electrochemical Water Oxidation and Oxygen Reduction. J. Am. Chem. Soc. 2013, 135, 8525−8534. (39) Shevchenko, D.; Anderlund, M. F.; Styring, S.; Dau, H.; Zaharieva, I.; Thapper, A. Water Oxidation by Manganese Oxides Formed from Tetranuclear Precursor Complexes: the Influence of Phosphate on Structure and Activity. Phys. Chem. Chem. Phys. 2014, 16, 11965−11975. (40) Marafatto, F. F.; Strader, M. L.; Gonzalez-Holguera, J.; Schwartzberg, A.; Gilbert, B.; Peña, J. Rate and Mechanism of the Photoreduction of Birnessite (MnO2) Nanosheets. Proc. Natl. Acad. Sci. U. S. A. 2015, 112, 4600−4605. (41) Birkner, N.; Nayeri, S.; Pashaei, B.; Najafpour, M. M.; Casey, W. H.; Navrotsky, A. Energetic Basis of Catalytic Activity of Layered Nanophase Calcium Manganese Oxides for Water Oxidation. Proc. Natl. Acad. Sci. U. S. A. 2013, 110, 8801−8806. (42) Krewald, V.; Retegan, M.; Cox, N.; Messinger, J.; Lubitz, W.; DeBeer, S.; Neese, F.; Pantazis, D. A. Metal Oxidation States in Biological Water Splitting. Chem. Sci. 2015, 6, 1676−1695.

27831

DOI: 10.1021/acsami.6b09984 ACS Appl. Mater. Interfaces 2016, 8, 27825−27831