Synthesis of New Mixed-Metal Ammonium Vanadates: Cation Order

Article pubs.acs.org/crystal

Synthesis of New Mixed-Metal Ammonium Vanadates: Cation Order versus Disorder, and Optical and Photocatalytic Properties Lan Luo, Erkang Ou, Tatyana I. Smirnova, and Paul A. Maggard* Department of Chemistry, North Carolina State University, Raleigh, North Carolina 27695-8204, United States S Supporting Information *

ABSTRACT: Two new ammonium vanadate hydrates, i.e., M3(H2O)2V8O24·2NH4 (M = Mn and Co, I and II, respectively) were synthesized using hydrothermal reaction conditions, and their structures were determined by single crystal X-ray diffraction [I: P2/ m (No. 10), Z = 1, a = 8.2011(2) Å, b = 3.5207(1) Å, c = 9.9129(3) Å, β = 110.987(2)°; II: C2/m (No. 12), Z = 2, a = 19.4594(6) Å, b = 6.7554(2) Å, c = 8.4747(3) Å, β = 112.098(2)°]. Interestingly, the two structures are homeotypic, with the structure of I exhibiting an uncommon type of structural disorder between locally-bridging Mn(H2O)22+ (i.e., part of the oxide framework) and nonbridging NH4+ cations over the same site (1:2 ratio), wherein two NH4+ ions occupy the same site as the two H2O molecules when Mn(II) is vacant. The amount of Mn(II) in the formula of I was determined by a combination of techniques, including electron paramagnetic resonance, while the relative amounts of NH4+/H2O in its structure were determined by combined thermogravimetric-mass spectrometry analyses as well as confirmed by infrared spectroscopy. In contrast, this site disorder is absent in the crystal structure of II, which contains a fully ordered arrangement of locally-bridging Co(H2O)22+ and NH4+ cations that alternate down its c-axis within a larger superstructure related to I by (a → c, b → 2b, c → 2a). Within both structures, the respective Mn2+/Co2+ cations bridge to neighboring edge-sharing chains of distorted VO5 square pyramids, forming a three-dimensional network that contains channels of H2O and NH4+ molecules. Hydrogen bonding distances in I are significantly longer and weaker than in II and leading to the disordered structure of I. Both show the loss of all H2O and NH4+ molecules, by ∼300 °C for I and a slightly higher ∼325 °C for II, in each case yielding V2O5 and MV2O6 (M = Co or Ni) as the final products. Both I and II exhibit visiblelight bandgap sizes of ∼1.55 and ∼1.77 eV, respectively, owing to low-energy metal-to-metal electronic transitions. Further, I shows a temperature-dependent photocatalytic activity (at 40 °C) for the production of hydrogen from the reduction of water under irradiation by UV−vis or only visible light at respective rates of ∼314 μmol H2 g−1 h−1 and ∼54 μmol H2 g−1 h−1 (irradiant power density of ∼1.0 W/cm2). Thus, these first two known ammonium vanadate hydrates provide new insights into the structural driving forces for ordered versus disordered structures as well as into their resulting physical properties. different atoms are distributed over symmetry-equivalent sites.6 However, some modes of disorder are relatively uncommon, such as arising from a disordered arrangement of covalently bonded molecules and metal cations around the same crystallographic sites. This type of disorder is described herein, wherein Mn(H2O)22+ and NH4+ cations are disordered over neighboring related sites within a microporous structure, in contrast to the fully ordered structural analogue comprised of Co2+ cations. From another perspective, research into metal oxides for photocatalytic water splitting has drawn intense interest as a route to solar-driven fuels production.7,8 While many different metal-oxide photocatalysts have been found to be active for hydrogen production,9,10 their activity has mainly been limited

1. INTRODUCTION Crystalline materials can exhibit different types of structural disorder that result in key changes to their symmetries and physical properties, providing new insights into structure− property relationships. For example, the cation disorder in Zn3Al6(PO4)12·4tris(2-aminoethyl)amine·17H2O results in a lowering of its space group symmetry from body centered to primitive cubic;1 the Li(Ni0.5Mn0.5)O2 compound with an intralayer cation disorder exhibits high rate capacity as a potential lithium battery cathode material;2 and the [NH4][Zn(HCOO)3] compound shows a disorder−order transition of the NH4+ cations that yields an unusual paraelectricferroelectric phase transition.3 Disorder in crystals can be induced, for example, by the disorder of guest molecules in porous structures,4 by flexible organic ligands,5 and by mixedcation site substitutions. These types of disorders can occur, for example, when one or more atoms are distributed over different crystallographic sites within the structure, or when two or more © XXXX American Chemical Society

Received: June 4, 2016 Revised: August 31, 2016

A

DOI: 10.1021/acs.cgd.6b00851 Cryst. Growth Des. XXXX, XXX, XXX−XXX

Crystal Growth & Design

Article

temperature programs: For I: 130 °C for 72 h; II: 165 °C for 10 h. These were ramped to 155 °C in 48 h, and then ramped to 140 °C in 10 h. All reactions were subsequently slowly cooled to room temperature at a rate of 6 °C/h. The products were washed with deionized water, collected, and dried at 60 °C overnight. Each compound could be obtained in high phase-purity (>95%) according to powder X-ray diffraction (PXRD) data (Figure S1 in Supporting Information). All percent yields were calculated based on the amount of vanadium loaded in each reaction. Compound I was synthesized by mixing 0.0367 g of Mn(NO3)2· 4H2O, 0.0400 g of NH4VO3, and ∼1 mL of H2O at a molar ratio of 1:2:∼300. Small brown needle-like crystals were obtained in a yield of ∼70%. Compound II was synthesized by mixing 0.0447 g of Co(NO3)2·6H2O, 0.0363 g of NH4VO3, and ∼0.5 mL of H2O in a molar ratio of 1:2:∼150. Brown needle-like crystals were obtained along with a small amount of yellow-colored fine powder, and which were separated by dispersing the product in ∼30 mL of water and sonicating for 20−30 s. The brown crystals settled to the bottom within ∼1−2 min, and the water was decanted containing the yellowcolored powder. This separation process was repeated three times until a high purity was achieved, as judged by PXRD data and visually under a microscope. The thin brown needle-like crystals of II were obtained in a yield of ∼34%. 2.3. Structural Characterization. Single-crystal data sets for I and II were collected on a Bruker APEX-II CCD diffractometer using graphite-monochromatized Mo Kα radiation (λ = 0.71073 Å) from a sealed tube source at 296 K. The initial unit cell determination and data reduction were performed within the Bruker SAINT program.23 Both structures were solved by direct methods, and the refinements were performed by full matrix least-squares methods within the software program SHELXS-97.24 Hydrogen atoms were placed in idealized positions and were fixed to ride on the parent nitrogen or oxygen atoms in I, and were allowed to freely refine in II. Selected crystallographic data and structure refinement parameters for each are given in Table 1. Selected interatomic distances, bond angles, and the

to ultraviolet-light owing to their large bandgap sizes.10 The discovery of new visible-light-driven photocatalysts has been one of the most challenging research topics over the past decade.11 Recently, mixed-metal vanadates have attracted considerable attention owing to their smaller bandgap sizes and visible-light photocatalytic properties, e.g., as found for BiVO4, Ag3VO4, and InVO4, which are active for dye degradation and water splitting under visible light.12−14 In each of these cases, the bandgap transitions are attributed to a metal-to-metal charge transfer transition, i.e., between the filled Bi 6s2 and empty V 3d0 orbitals in BiVO4.15 Importantly, these materials demonstrate the ability to harness visible light in order to drive photocatalytic reactions for either water purification or hydrogen production.16−18 Recently, research in the Maggard group has demonstrated that the incorporation of high-spin Mn(II) cations into vanadate structures results in a significant red-shifting of their bandgap sizes, e.g., by up to ∼0.6 eV deeper into the visible-light wavelengths.19 For example, a family of manganese-vanadate organic hybrids was found to show temperature-dependent visible-light photocatalysis for water reduction, including Mn(terpy)V2O6, Mn(ophen)(H2O)V2O6, Mn(4,4-bpy)V2O6, and Mn2(bipym)V4O11, etc.18,20,21 The first example of this family also included Mn(2,2′bpy)V4O10(2,2′-bpy), wherein the coordination of the ligand to both the vanadium and manganese cations was postulated to facilitate enhanced charge separation within its structure.18 However, investigations into the optical properties and photocatalytic activities of mixed-metal vanadates are still few in number. In this work, two new ammonium vanadate hydrates: Mn3(H2O)2V8O24·2NH4 (I) and Co3(H2O)2V8O24·2NH4 (II) have been synthesized via hydrothermal methods. As the first ammonium manganese vanadate hydrate reported, I exhibits a rare structural disorder which consists of locally-bridging Mn(H2O)22+ and nonbridging NH4+ cations distributed over the same sites within the structure. Its structure and composition have been characterized by IR, EDS, and quantitative EPR spectroscopic analyses for the amount of Mn(II) cation within the structure. Additionally, in situ TGAMS measurements were utilized to characterize and confirm the amounts of H2O and NH4+ within the structure of I.22 Its structure and composition are contrasted with the highly ordered structure of the cobalt analogue in compound II, wherein the Co(H2O)22+ and two NH4+ cations alternate within its crystalline structure. The superstructure of II results in a unit cell that is four times larger than for I, with a change in the space group symmetry from P2/m to C2/m. The syntheses, structures, thermal stabilities, and optical and photocatalytic activities for I and II are described.

Table 1. Selected Crystal Data and Structure Refinement Details for I and II

2. EXPERIMENTAL SECTION

compound

I

II

formula crystal system space group, Z temperature, K a, Å b, Å c, Å β, ° V, Å3 ρcalc, g/cm3 μ, mm−1 total reflections, Rint data/restraints/parameters final R1, wR2a [I > 2σ(I)] Δρmax/Δρmin, e/Å3

Mn1.5V4O13NH6 monoclinic P2/m, 1 296.2 8.2011(2) 3.5207(1) 9.9129(3) 110.987(2) 267.23(1) 3.16 5.15 7789, 0.0428 1530/0/64 0.0362, 0.0936 1.06/−1.79

Co3V8O26N2H12 monoclinic C2/m, 2 100.1 19.4594(6) 6.7554(2) 8.4747(3) 112.098(2) 1032.21(6) 3.35 5.91 13996, 0.0398 2323/0/123 0.0241, 0.0557 1.06/−0.83

a R1 = Σ(|Fo − Fc|)/ΣFo; wR2 = [Σ(w(Fo2 − Fc 2)2)/(Σ(Fo2)2) ]1/2; w = σF−2.

2.1. Materials. All reagents were used as supplied by the manufacturer without further purification, including Mn(NO3)2· 4H2O (99.98%, Alfa Aesar), Co(NO3)2·6H2O (98.0%, Alfa Aesar), and NH4VO3 (99.99%, Aldrich). A reagent amount of deionized water was also used in each reaction. 2.2. Synthetic Procedures. The M3(H2O)2V8O24·2NH4 compounds (I, M = Mn; II, M = Co) were prepared using hydrothermal techniques by loading stoichiometric amounts of reactants and deionized water into polytetrafluoroethylene Teflon pouches (3″ × 3″), and were then heat sealed in air. These pouches were placed into 125 mL polytetrafluoroethylene-lined stainless-steel autoclaves that were backfilled with ∼40 mL of deionized water before closing. The reaction vessel was heated in a convection oven with the following

results of bond valence sum calculations are listed in Tables S1 and S2 in the Supporting Information. The crystalline purities of the bulk synthetic products were characterized by high-resolution PXRD on a RIGAKU R-Axis Spider powder X-ray diffractometer (graphite monochromatized Cu Kα radiation) at room temperature. PXRD patterns were scanned with a step size of 0.02° over the 2θ angular range from 4° to 100° and dwell times of 4 s for each step. 2.4. Spectroscopic and Thermal Characterization. Midinfrared (400−4000 cm−1) spectra were measured on an IR-Prestige B



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Figure 1. Polyhedral views of the ordered structure of II (left) and disordered structure of I (right) down the b axis in both; the yellow polyhedra are centered by vanadium, blue polyhedra are centered by cobalt, red atoms are oxygen, and white atoms are hydrogen. Local hydrogen bonds are drawn as dashed lines. 17) and NH2+ (mass 16) with a relative abundance of 0.675:1.28 A control experiment on NH4VO3 was performed. The relative molar amounts H2O and NH3 were quantitatively determined from calculating the signal area of each mass based on the relative abundances. 2.5. Photocatalytic Activity. Photocatalytic activities for hydrogen production were measured by suspending a weighed amount (∼10 mg) of each powdered sample within an outer-irradiation range quartz reaction cell which was filled with ∼45 mL of a 20% aqueous methanol solution. In order to remove any trapped gases on the particles’ surfaces, the particle suspension was first stirred in the dark for 30 min with constant nitrogen purging, and then sonicated in degas mode for 90 min. The reaction cell was then irradiated under an 800 W high pressure Xe arc lamp (Newport Oriel 6271, focused through a shutter window) which was equipped with an external fan, an IR water filter, and visible-light cutoff filter (λ > 420 nm). The reaction temperature was controlled using an external heater and measured every 15 min. The reaction temperature was slowly increased until a measurable amount of hydrogen was produced within an initial reaction time of 15 min. Measurable rates for photocatalytic activities were found starting at 40 °C for I. The solution was irradiated with continuous stirring under visible light at an irradiant power density of ∼1 W/cm2. The reaction cell was connected to an L-shaped horizontal quartz tube in order to collect the evolved gas. The amount of the gases produced was determined volumetrically every 30 min. The collected gas was manually injected into a gas chromatograph (SRI MG #2; thermal conductivity detector) to identify the generated gases and to check for a constant molar ratio over time. Only hydrogen gas was detected by gas chromatography. Carbon dioxide was not detected in the gases products, as is most commonly observed when methanol is used as the sacrificial hole scavenger (i.e., CH3OH + H2O → 3H2 + CO2). Tollen’s tests were used to confirm aldehyde as the primary oxidation product for each of the photocatalytic reactions.

21 Shimadzu Fourier transform infrared spectrophotometer (FTIR) equipped with a GladiATR accessory. Approximately 10 mg of highpurity powder was loaded onto the sample stage, and the data were plotted as transmittance versus wavenumber. UV−vis diffuse reflectance spectra (DRS) were measured on a Shimadzu UV-3600/ 3100 UV−vis-NIR spectrophotometer equipped with an integrating sphere. A pressed barium sulfate powder tablet was used as the reference. The data were plotted as the remission function F(R∞) = (1 − R∞)2/(2R∞) versus energy, where R∞ is diffuse reflectance according to the Kubelka−Monk theory of diffuse reflectance.25 The optical bandgap sizes were estimated with the formula Eg (eV) = 1240/λg (nm), where λg is extrapolated from the linear rising part in the spectra.26 Continuous wave (CW) X-band EPR spectra were collected using a Bruker Elexsys E500 spectrometer at room temperature at microwave frequencies of approximately 9.5 GHz (Xband). Each solid sample was weighed (∼0.45 mg) and placed at the bottom of a 3 mm quartz EPR tube for measurements. Solutions were drawn into a glass capillary with i.d. = 0.0276 in. (Drummond Scientific Company, Broomall, PA) to occupy approximately 5 cm of the length of the capillary, sealed, and placed into standard 3 mm quartz EPR tube for measurements. CW EPR spectra were digitized to 2048 data points per spectrum. Typical spectrometer settings were as follows: microwave power of 2 mW (20 dB attenuation), field modulation frequency of 100 kHz, and field modulation amplitude of 10 G. Energy dispersive spectroscopy (Oxford Instrument, INCA PentaFET−x3, built in a Hitachi S-4700 cold cathode field-emission SEM) measurements were taken on each sample at an operating voltage of 20 kV. All samples were coated with an ∼3 nm Au/Pd layer using a Cressington 108 auto sputter coater prior to the measurement to prevent charging of the samples. Thermogravimetric analyses were taken on a TA Instruments TGA Q50 by loading 5 mg onto a platinum pan. The pan was equilibrated and tared at room temperature and heated to 600 °C at a rate of 5 °C/ min under flowing N2 or air. The post-heated TGA residues were characterized by PXRD, as described above. Specific surface areas were measured on a Quantachrome ChemBET Pulsar TPR/TPD. The samples were degassed by heating to 160 °C under flowing N2 for 3 h, and then cooled with liquid nitrogen and analyzed using a 30% He/N2 gas mixture. The gaseous products evolving from the thermal decomposition of I were characterized by combined TGA-MS techniques (TA Instruments, Discovery TGA coupled with a Discovery quadrupole mass spectrometer). The mass scans in the TGA-MS experiments were done using an ionization energy of 40 eV and a heating rate of the sample of 10 °C/min. The quantitative analysis of H2O and NH3 can be determined by the possibility of ionization (ionization cross section). At 40 eV, the major ionization product of H2O through electron impact are H2O+ (mass 18) and OH+ (mass 17) with a relative abundance of 1:0.251;27 the major ionization products of NH3 through electron impact are NH3+ (mass

3. RESULTS AND DISCUSSION 3.1. Structural Description. 3.1.1. Mn3(H2O)2V8O24·2NH4 (I). The structure of I consists of edge-sharing vanadate and manganese oxide chains that are oriented down the b axis, and are linked through vertical oxygen atoms, as shown in Figures 1 (right) and 2 (bottom). There are two symmetry-unique vanadium atoms (V1, V2) and two types of manganese atoms (Mn1, Mn2). The vanadate zigzag chains along the b axis are constructed from edge-sharing VO5 square pyramids connected via basal oxygen atoms (O3, O4). For each VO5 square pyramid, the V−O basal bonds [V−O at 1.662(2)−2.131(3) Å] are significantly longer than the V−O apical bonds [V−O at 1.611(3)−1.651(3) Å], consistent with distances previously reported in vanadate structures.33 The two parallel vanadate C



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other manganese vanadates, e.g., as for MnV2O6, Mn2V2O7, etc.36,37 Within the structure, the coordinating water molecules also interrupt the [-O−Mn−O−V−O-] bridging of the metal− oxygen network. The coordination geometry of Mn1 is also altered, with the VO5 pyramids arranged in a trans fashion as compared to the cis coordination for Mn2. The hydrogen bonding distances between the donor and acceptor are all at ∼2.85 Å, Figure 1 (dashed lines), i.e., for N−H···O or O−H··· O between the NH4+/H2O molecules and the apical oxygen atoms of VO5. Hydrogen bonding at this donor−acceptor distance has been classified as a moderate and a mostly electrostatic interaction.35 The randomized occupation of this site by either locally-bridging Mn(H2O)22+ or two NH4+ has no impact on this hydrogen-bonding distance as a result of the regularly repeating apical O−O distances of 3.52 Å. This distance is determined by the b axis length and is labeled with a doubled arrow in Figure 2. This compares to the structure of II, wherein alternately longer and shorter O − O distances repeat down the b axis in order to optimize the hydrogen bonding interactions via an ordered structural arrangement (described below). 3.1.2. Co3(H2O)2V8O24·2NH4 (II). As divalent transition metals, the Mn(II) and Co(II) cations exhibit similarities in their preferred coordination geometries and environments.38,39 While many Mn(II)-containing compounds have been found to be isostructural with the Co(II)-containing analogues, many others can reveal key structural differences.40−42 Thus, synthetic efforts were aimed at the synthesis of the Co(II) analogue and thereby possibly gaining further insight into the origins of the structural disorder in I. Compound II, the Co(II)-containing analogue of I, can be obtained at relatively higher temperatures and using less solvent water during the reaction. Structural characterization of II shows it is an ordered version of I that crystallizes in the space group C2/m with a unit cell that has quadrupled in size. As shown in Figures 1 and 2, the structure of II consists of edge-sharing vanadate chains linked through vertical oxygen atoms from edge-sharing cobalt oxide chains. Within these chains, the two types of cobalt atoms (Co1 and Co2) are coordinated in distorted octahedral geometries, while two of the vanadium atoms (V1, V3) are coordinated in tetrahedral geometries and the third vanadium in a square pyramidal geometry (V2). Neighboring VO5 square pyramids and VO4 tetrahedra form ladder-like chains along the b-axis direction via corner-sharing oxygen atoms (O3 and O4), shown in Figure 2. Two of the vanadate chains and two CoO6 octahedral chains are bridged in parallel and linked via the basal oxygen atoms (O2 and O5) to form a channel along the b-axis. Each channel is bridged via the opposite vertices of the MnO6 octahedra to form layers within the (0 0 1) plane. Each layer is linked through the Co(H2O)2O4 octahedra via apical oxygen atoms (O6) from the VO5 pyramids, forming a three-dimensional network as shown in Figure 2. In the VO5 square pyramid, the V−O basal bonds [V−O at 1.709(2)−2.075(2) Å; typical V−O basal bonds are ∼1.80−2.01 Å]33 are significantly longer than the V−O apical bonds [V2−O2 at 1.620(2) Å], as described for I. The V−O bonds in the VO4 tetrahedra are similar to the V− O apical bonds in the VO5 pyramids [V−O at 1.612(2)− 1.851(1) Å]. The Co(H2O)2O4 octahedral environment consists of two water molecules [Co−OH2 at 2.072(1) Å; a typical Co−OH2 bond is 2.139(1) Å]43,44 and four bridging oxygen atoms [Co−O at 2.056(3)−2.072(1) Å] that are bridged to the VO5 square pyramids and VO4 tetrahedra. The

Figure 2. Polyhedral views of the disordered structure of I (bottom) and ordered structure of II (top) down the a axis and c axis, respectively; yellow polyhedra are centered by vanadium, blue polyhedra are centered respectively by cobalt or manganese, red atoms are oxygen, and white atoms are hydrogen. Hydrogen bonds are drawn as dashed lines, and the double arrows labeled (1) and (2) indicate key O−O distances, described in text.

chains are oxo-bridged by MnO6 octahedral chains of Mn2 via the remaining basal oxygen atoms (O2 and O5), forming channels along the b-axis and oxide layers within the (0 0 1) plane. The oxide layers are further linked through Mn1 into a three-dimensional network, as shown in Figure 1. Each Mn1 cation is coordinated to four apical oxygen atoms (O6) from four different VO5 pyramids in the equatorial plane [Mn−O at 2.282(2) Å] and to two water molecules in the axial direction [Mn−OH2 at 2.122(3) Å; a typical Mn−OH2 bond is 2.21(1) Å].34 The Mn−O bonds in MnO6 octahedra (Mn2) are shorter than those in the equatorial plane of Mn1 in Mn(H2O)2O4 [Mn−O at 2.053(3)−2.235(2) Å]. Interestingly, I exhibits a rare type of disorder of the locallybridging Mn(H2O)22+ cations (i.e., that are part of the oxide framework) with two nonbridging NH4+ cations, wherein each are distributed over neighboring sites at 50% occupancy. This isoelectronic substitution is necessary for charge balancing and stabilizing the Mn2+ vacancies in the structure. As a result, relatively open channels are formed along the b axis in the structure, making the compound less dense as compared to D



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Co−O distances within the CoO6 octahedra [Co−O at 1.979(3)−2.179(1) Å] are at relatively shorter distances than for the Mn−O distances of the MnO6 octahedra in I. Notably, the crystallographic sites for locally-bridging Co(H2O)22+ (i.e., that are part of the oxide framework) and NH4+ cations are both fully occupied, i.e., at sites with half the Wyckoff multiplicity as compared to I, and alternate in an ordered arrangement down the chains in the b-axis direction. As compared to the structure of II, the averaged b-axis length is a shorter 3.375 Å (i.e., b/2), or by ∼0.15 Å, and yields shorter hydrogen bonding distances in the structure. This results from the relatively smaller atomic radius of Co(II) and the shorter Co−O distances as compared to Mn(II). However, the O−O distances down the b axis alternate between longer and shorter at ∼3.88 and 2.87 Å, labeled in Figure 2, in order to accommodate the hydrogen bonding environments of NH4+ and Co(H2O)22+, respectively. As a result, the V−O−V angles are bent at a more acute ∼134° in II, as compared to the same V−O−V angles in I at ∼142°. The NH4+ cations are hydrogen bonded to the O atoms at donor−acceptor, i.e., N−H···O, distances of ∼2.9 to ∼3.1 Å ( × 5). The water molecules are hydrogen bonded at shorter distances, i.e., O−H···O, at ∼2.8 Å (× 2). Thus, a significantly greater amount of hydrogen bonding is achieved in the structure of II, as a result of a shorter b-axis length that facilitates the cation-ordered structure. 3.2. Composition Characterization. As described in the preceding section, the structural refinements of I indicate that one of the manganese atoms (Mn1) is located in a halfoccupied site within the unit cell. Each Mn(II) cation is coordinated to two water molecules, e.g., as found in the Co(II)-containing structure of II. Interestingly, while Mn1 is 50% occupied, its coordinating axial sites are found to be fully occupied. This unusual observation arises because when the site for Mn(II) is vacant, its two coordinating water molecules are replaced with two NH4+ cations at the same crystallographic site, i.e., as more clearly observed within the structure of II. Unsuccessful attempts were made to refine the structure of I in alternate space groups and unit cell settings, in order to eliminate the structural disorder and find an ordered arrangement. However, no ordered structural versions could be successfully refined. Further, it is difficult to distinguish nitrogen and oxygen using the single crystal X-ray refinements, thus raising the possibility of a deviation of the chemical composition of I from that found in the ordered structure of II, i.e., a mixture of NH3, NH4+, H2O, or H3O+, or a combination thereof. Thus, further spectroscopic and elemental characterizations were performed on I in order to more completely characterize its composition and disordered structure. As shown in Figure 3, the infrared spectrum of I exhibits broad absorption bands in the 400 to 1200 cm−1 range, which are attributed to both V−O and Mn−O vibrations. The O−H groups from the coordinated water molecules are present as broad peaks centered roughly around 3300−3500 cm−1. Similar to II, compound I exhibits two sharp peaks at 1579 and 1400 cm−1, corresponding to the bending mode of water molecules and NH4+ respectively, which are in good agreement with the functional groups in the IR spectra of H2O and NH4VO3, as compared in Figure 3.29−32 Characteristic peaks of neither the hydronium ion at 1715 and 2150 cm−1, nor coordinating NH3 molecules at 1205 cm−1 were observed in the spectra.29−32 TGA-MS methods were utilized to quantify the relative amounts of water and ammonium cations in the structure of I, as shown in Figure 6. As a reference, the TGA-MS of

Figure 3. IR spectra of I (black), II (green), NH4VO3 (red), and H2O (blue).

NH4VO3 was performed in order to calibrate the signals arising from NH3 and H2O. A quantitative calculation of NH3 and H2O evolved from I was obtained from the signal integrations of the corresponding mass numbers (i.e., 16, 17, and 18 AMU). The calculated ratio of H2O and NH3 evolved from NH4VO3 based on the peak area was 0.451 (cald: 0.50, error: 9.7%), while the ratio of H2O and NH3 in I was found to be 1.46 (cald: 1.50, error: 2.39%). Therefore, the IR and TGA-MS results are consistent with the 1:1 molar ratio of H2O and NH3 in the structure of I. The X-band EPR results on I are provided in Figure S2 in the Supporting Information. The fine powder sample of I exhibits a single line spectra with a corresponding peak-to-peak width of 255 G, and no corresponding resolved anisotropy of hyperfine features. The effective g value for compound I was calculated to be 2.005(1), in agreement with the high-spin Mn(II) ions, (g = 2.000) in a distorted octahedral coordination environment within its structure.47 The observed symmetric EPR signal with no resolved g-anisotropy or hyperfine features is a result of strong Heisenberg spin exchange interaction. The two manganese centers, i.e., Mn1 and Mn2, have different coordination environments, and as a result, are magnetically nonequivalent and may have different hyperfine and g-tensors. However, the distance between the nearest Mn1 and Mn2 centers is sufficiently short, ∼5.47 Å, to ensure a strong spin exchange interaction resulting in a featureless single-line EPR signal. Bond valence calculations for the manganese sites in each compound are consistent with the Mn(II), Co(II), and V(V) oxidation states, as listed in Tables S1 and S2 for I and II in the Supporting Information. In order to characterize the amount of Mn and V in I, energy dispersive spectroscopic data were taken on its powder, as shown in Figure 4 for six different sites within the crystalline mixture. Results of the corresponding elemental analyses are in agreement with the atomic ratio for Mn:V of 3:8 in the formula (Mn:V ratio: exp. 0.385, cald. 0.375). In order to further quantify the Mn content in the bulk powder of I, solution-based EPR measurements were performed using prepared solutions of Mn(CH3COO)2 as a reference in order to calibrate the observed signals to the amount of Mn(II). Four solutions were prepared for the quantitative EPR analysis using 1 mol/L HCl solution as the solvent: (a) 3.33 mmol/L solution using 3.408 mg of I; (b) 1.67 mmol/L solution using 1.704 of I; (c) 10 mmol/L Mn(CH 3 COO) 2 solution: 2.451 mg of MnE



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Figure 4. Images from scanning electron microscopy (top) and energy dispersive spectra (middle) from two sample areas for the powder of I. The integrated atomic percentages from six different trials are given in the table (bottom).

(CH3COO)2·4H2O; (d) 5 mmol/L Mn(CH3COO)2 solution: 1.225 mg of Mn(CH3COO)2·4H2O. The amount of Mn(II) in solutions a and b can be calculated from the double integration of EPR spectra by calibrating the signals to that measured for reference solutions c and d. The results show that solution a (3.33 mmol/L of I) exhibits the same EPR signal as solution c (10 mmol/L Mn2+), and solution b (1.67 mmol/L of I) shows the same EPR signal with solution d (5 mmol/L Mn2+), shown in Figure 5. These results indicate that the amount of Mn(II) in I is 3 mol per formula, and which is consistent with the crystallographic refinements of its compositions described above, i.e., for Mn3(H2O)2V8O24·2NH4. 3.3. Thermal Stability. The low-temperature removal of small molecules (NH3, H2O etc.) from metal oxides represents a potentially promising route to prepare new microporous and/ or nanostructured materials.45,46 The thermal stabilities and decomposition products of I and II were investigated by heating their powders under flowing nitrogen at a ramping rate of 5 °C per minute. The thermogravimetric results of the weight losses versus temperature are shown in Figure 8. The post-TGA products were characterized by PXRD and specific surface area analysis. Compounds I and II show a relatively comparable range of thermal stabilities and exhibit the loss of structural water and

Figure 5. EPR spectra of dissolved I with concentrations of 6.66 mmol/L (red line) and 3.33 mmol/L (blue cross), and Mn(CH3COO)2 solution of 5 mmol/L (green line) and 10 mmol/L (black dot). The solvent is 1 mol/L HCl.

ammonia with heating to similar temperature ranges. Shown in Figure 8, both I and II exhibit a single weight-loss step for the removal of structural water and the ammonium cations. F



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Figure 7. Tauc plots of [F(R)hν]n vs [hν] for the direct and indirect band gap transitions, n = 2 (right) and n = 1/2 (left), respectively, of I (black) and II (blue).

shows a higher-energy bandgap size of ∼1.77 eV. Both show light absorption at visible-light wavelengths, with the smaller band gap of I arising from the higher energy Mn d orbitals. The heterogeneous photocatalytic activity of I for hydrogen production was investigated as a particle suspension within a 20% aqueous methanol solution that was irradiated by UV and visible light, or only visible light. Compound I was surface impregnated with a standard 1% (by weight) Pt surface cocatalyst, as described previously.18,48 PXRD was used to confirm the bulk stability of I under these photocatalysis testing conditions, as provided in the Supporting Information. Figure 6. TGA-MS curves for I (top) and NH4VO3 II (bottom) under argon; black: plotted as weight (%) versus temperature (°C), blue: 16 AMU, green: 17 AMU, red: 18 AMU.

Systematic variation of the heating rate and temperature could not be used to achieve separate weight-loss steps for the water molecules and the ammonium cations. At the heating rate of 5 °C/min, both the water and ammonium cations are removed starting at ∼200 to 320 °C (exp. 8.57%, calcd. 8.56%) in I. In II, the water and ammonium cations are observed to be removed at ∼165−330 °C (exp. 8.48%, calcd. 8.46%). After the removal of water and ammonia, I and II are found to decompose into MV2O6 (M = Mn for I (ICSD #40850); M = Co for II (ICSD #40849)) and V2O5 (ICSD #60767), as confirmed by PXRD given in the Supporting Information. The net decomposition reactions can be represented by the following equations: Mn3(H2O)2V8O24·2NH4 → 3MnV2O6 + V2O5 + 3H2O + 2NH3 for I, and Co3(H2O)2V8O24·2NH4 → 3CoV2O6 + V2O5 + 3H2O + 2NH3 for II. 3.4. Optical and Photocatalytic Properties. Optical UV−vis DRS measurements were taken for I and II, in order to determine the bandgap sizes and the onset of absorption for photocatalysis measurements. Typically, simple vanadate structures exhibit bandgap sizes of ∼2.0 eV to ∼2.4 eV. However, the incorporation of the Mn(II) and Co(II) cations within their structures can cause a redshift of their optical absorption owing to the introduction of partially-filled dorbitals located within their band gaps. Optical absorptions arising from the (spin-forbidden) d-d transitions on the Mn(II) and Co(II) cations were not observable in any of the spectra. Bandgap sizes (indirect) were found in the visible region of energies for each compound, as shown in the Tauc plots in Figure 7. I exhibits a lower-energy band gap of ∼1.55 eV, and II

Figure 8. Thermogravimetric analysis of compound I (red) and II (black) under nitrogen with 5 °C/min ramping up, plotted as weight (%) versus temperature (°C).

At room temperature, i.e., ∼ 25 °C, I was found to be photocatalytically inactive. However, measurable photocatalytic activities begin to emerge at an elevated temperature of ∼40 °C. Approximate turnover numbers for the amount of hydrogen produced per surface site were calculated based on the measured specific surface areas and the average density of surface sites calculated for various surface terminations of the crystal structures, as described previously.18 Time course plots of the turnover numbers are shown in Figure 9. At 40 °C, compound I exhibits a high initial photocatalytic activity of ∼314 μmol H2 g−1·h−1, with a TOF of 87.5 h−1 under UV and G



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4. CONCLUSIONS The first ammonia manganese vanadate hydrate (Mn 3 (H 2 O) 2 V 8 O 24 ·2NH 4 ; I) and its cobalt analogue (Co3(H2O)2V8O24·2NH4; II) have been synthesized hydrothermally, and their crystalline structures, thermal stabilities, optical properties, and photocatalytic activities have been investigated. Both I and II exhibit the same structural connectivity, with the exception that I exhibits an uncommon structural disorder consisting of the statistical distribution of Mn(H2O)2+ and NH4+ over related neighboring crystallographic sites. By contrast, compound II consists of a fully ordered superstructure containing alternating Co(H2O)2+ and NH4+ cations down the b-axis direction. The optical bandgap sizes were measured to be ∼1.55 eV for I and ∼1.77 for II. Notably, I is photocatalytically active for hydrogen production in aqueous solutions under UV and visible light, as well as under only visible light, at rates of ∼314 μmol H2 g−1 h−1 with a TOF of ∼87.5 h−1, and ∼54 μmol H2 g−1 h−1 with a TOF of ∼12 h−1, respectively. Thus, these results represent the first observed type of structural disorder between ammonia and metal cations, as well as shedding new light on the application of manganese vanadates as visible-light photocatalysts.

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Figure 9. Photocatalytic hydrogen production rates per surface formula of I under UV light at 250−800 nm (top) and visible light at 420−800 nm (bottom). Testing conditions for hydrogen production included catalyst (∼10 mg) with 1 wt % Pt as a surface cocatalyst suspended in a 20% methanol solution (∼45 mL), at a radiant power density of 1 W/cm2 at 40 °C.

ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.cgd.6b00851. Tables of selected interatomic distances for I and II, powder X-ray diffraction data of the as-prepared products, after heating to 500 °C under nitrogen, and after photocatalytic testing, and infrared and electron paramagnetic resonance spectra (PDF)

visible light in the first 2 h, before gradually slowing after 4 h. However, its activity is fully recovered by washing and resuspending the catalyst powder in a fresh 20% methanol aqueous solution, as shown in Figure 9. Under visible light, the initial photocatalytic activity of I is ∼54 μmol H2 g−1·h−1, with a TOF of ∼12 h−1. The increased photocatalytic rates at higher temperatures likely result from the faster adsorption/ desorption rates at the surfaces of metal oxides.21 These rates are comparable to photocatalytic rates for hydrogen production for related manganese-vanadate hybrids, i.e., which span a range of ∼100 to 670 μmol H2 g−1·h−1.18,21 The highest rates have been observed with increased temperature and with an increase in the dimensionality of the oxide/organic framework, as described for the structure of I above. Further investigations are necessary to probe the temperature-dependent mechanism of its photocatalytic activity. Notably, carbon dioxide was not detected as a reaction product in any of these reactions, as is commonly observed in photocatalytic reactions using methanol as a hole scavenger. All post-reaction solutions exhibited a positive Tollen’s test and a decrease of pH value, which indicated the partial oxidation of methanol to formic acid. These results are similar to photocatalytic reactions involving oxide/organic hybrids, e.g., manganese vanadate hybrids, Cu(I)- and Cu(II)-molybdate hybrids, and which were also recently shown to result in the formation of formic acid and formaldehyde, rather than the complete oxidation of methanol to carbon dioxide.21,49 Thus, inorganic manganese vanadates show promising photocatalytic rates for visible-light photocatalytic hydrogen production and yield similar reaction products as found for the mixed-metal hybrid systems.18,21

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AUTHOR INFORMATION

Corresponding Author

*Contact Email: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors acknowledge support of this research from the Department of Chemistry at North Carolina State University. T.S. and E.O. were supported by NSF-1508607. The EPR instrumentation at North Carolina State University is supported by Grants NIH-S10RR023614, NSF CHE0840501, and NCBC 2009-IDG-1015.

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Synthesis of New Mixed-Metal Ammonium Vanadates: Cation Order

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