Open-Framework Manganese(II) and Cobalt(II) Borophosphates with

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Open-Framework Manganese(II) and Cobalt(II) Borophosphates with Helical Chains: Structures, Magnetic, and Luminescent Properties Min Li,† Volodymyr Smetana,‡,§ Magdalena Wilk-Kozubek,‡,§,∥,⊥ Yaroslav Mudryk,‡ Tarek Alammar,‡,∥ Vitalij K. Pecharsky,‡,∥ and Anja-Verena Mudring*,‡,§,∥ †

Department of Chemistry, Iowa State University, Ames, Iowa 50011-3111, United States Ames Laboratory, U.S. Department of Energy, Iowa State University, Ames, Iowa 50011-3020, United States § Department of Materials and Environmental Chemistry, Stockholm University, Svante Arrhenius väg 16 C, 10691 Stockholm, Sweden ∥ Department of Materials Science and Engineering, Iowa State University, Ames, Iowa 50011-2300, United States ⊥ Department of Nanotechnology, Wrocław Research Centre EIT+, 147 Stabłowicka Street, 54-066 Wrocław, Poland ‡

S Supporting Information *

ABSTRACT: Two borophosphates, (NH4)1−2xM1+x(H2O)2(BP2O8)· yH2O with M = Mn (I) and Co (II), synthesized hydrothermally crystallize in enantiomorphous space groups P6522 and P6122 with a = 9.6559(3) and 9.501(3) Å, c = 15.7939(6) and 15.582(4) Å, and V = 1275.3(1) and 1218.2(8) Å3 for I and II, respectively. Both compounds feature helical chains composed of vertex-sharing tetrahedral PO4 and BO4 groups that are connected through O atoms to transition-metal cations, Mn2+ and Co2+, respectively. For the two crystallographically distinct transition-metal cation sites present in the structure, this results in octahedral coordination with different degrees of distortion from the ideal symmetry. The crystal-field parameters, calculated from the corresponding absorption spectra, indicate that Mn2+ and Co2+ ions are located in a weak octahedral-like crystal field and suggest that the Co−ligand interactions are more covalent than the Mn−ligand ones. Luminescence measurements at room temperature reveal an orange emission that red-shifts upon lowering of the temperature to 77 K for I, while II is not luminescent. The luminescence lifetimes of I are 33.4 μs at room temperature and 1.87 ms at 77 K. Both compounds are Curie−Weiss paramagnets with negative Weiss constants and effective magnetic moments expected for noninteracting Mn2+ and Co2+ cations but no clear long-range magnetic order above 2 K.

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INTRODUCTION Borophosphates (BPOs) are zeolitic materials that often feature open-framework structures with wide pores through the connection of metal polyhedral or organic species with oxygen-interlinked borate and phosphate anions. In comparison to aluminophosphates,1,2 borates may exhibit both planar Bψ3 and tetrahedral Bψ4 (ψ = O, OH) units, which opens up additional structural space for open-framework materials. In addition, various main-group and transition-metal cations have been successfully incorporated in BPOs, adding further to the structural complexity of this class of compounds.3−6 Moreover, organic cations have been explored as templates to guide the structure formation and ultimately control the pore size.4−10 The introduction of new synthetic techniques such as the ionothermal synthetic approach11−14 has allowed for the preparation of BPOs that are not attainable by other methods. As a result, a broad spectrum of anion-partial BPO structures, including oligomeric units,6,15,16 chains,17−19 ribbons,20−22 layers,23−25 and three-dimensional frameworks,26−28 has been reported to date. The structural diversity of BPOs also brings © XXXX American Chemical Society

along a variety of different functionalities. Examples include nonlinear-optical,29−31 luminescent,14,32 magnetic,33−36 and lithium- and sodium-ion-battery materials.37,38 For MeBPOs (metal−BPO systems), the following representatives have been reported: MIMII(H2O)2[BP2O8]·yH2O, with MI = Li, Na, K, Rb, Cs, and NH4, MII = Mg, Sc, Mn, Fe, Co, Ni, Cu, Zn, and Cd, and y = 0.2−1 for mono- and divalent cations,22,39−48 and MIII(H2O)2[BP2O8]·H2O, with MIII = Fe48 and In,49 for triply charged cations featuring a complex [BP2O83−] unit as the vertex-sharing BO4 and two PO4 tetrahedra (Figure S1). The connection of these units to extended chains results in the formation of helices that are either left- or right-handed. For charge balance, cations occupy the space between the helices. The helices can be considerably compressed or expanded depending on the different sizes and charges of embedded cations. An example for the tremendous structural flexibility of this structure type is [C4N3H16]Received: June 6, 2017

A

DOI: 10.1021/acs.inorgchem.7b01423 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

X-ray photoelectron microscopy (XPS) measurements were carried out at RT on a physical electronic 5500 multitechnique system with a standard aluminum source. The analysis spot size was 1 × 1 mm2. Samples were mounted on double-sided Scotch tape. The binding energies in the XPS spectra were calibrated against the C 1s signal (284.8 eV) corresponding to adventitious physisorbed carbon dioxide. Absorption spectra were recorded on powder samples using an Agilent Technologies Cary 5000 UV−vis−near-IR (NIR) spectrophotometer equipped with an internal diffuse-reflectance accessory (Praying Mantis, Harrick). To avoid signal saturation, a sample of compound II was ground with BaSO4 in the mass ratio of 1:1 prior to the measurement. Optical-grade BaSO4 was used as a reference. Corrected steady-state excitation and emission spectra were recorded on a Horiba Jobin Yvon Fluorolog 3-222 spectrofluorometer, equipped with a 450 W xenon arc lamp and a R928P photomultiplier-tube detector. The decay curve measurements were carried out on the same instrument operating in phosphorescence mode. In those measurements, a xenon flash lamp was used for sample excitation. Absorption spectra and luminescence measurements were conducted at RT and 77 K. Magnetic measurements of polycrystalline powder samples were performed using a superconducting quantum interference device (SQUID) magnetometer (model MPMS XL-7, Quantum Design, USA) in the temperature range from 2 to 300 K and in magnetic fields from 0 to 7 T.

[Zn3B3P6O24]·H2O, which contains large organic cations, forcing a stretch of the chains. 46 Another interesting representative is (NH4)0.75Fe(H2O)2[BP2O8]·0.25H2O, which contains both Fe2+ and Fe3+ and has been discussed as a cathode material for sodium-ion batteries.37 Moreover, the presence of two crystallographically different types of water in the structure of MIMII(H2O)2[BP2O8]·yH2O results in a more complex thermal behavior.43 However, despite the numerous reports of the preparation and structural chemistry, relatively little research is devoted to the properties of these crystalline compounds compared to other zeolitic open-framework solids such as AlPOs (Al−P−O systems),50−52 GaPOs (Ga−P−O systems),53−55 and MeAlPOs (metal−Al−P−O systems),56−58 which have extensively been researched for applications ranging from microelectronics to medical diagnosis.59,60 Thus, property investigations on BPOs, especially for the subgroup of chiral of BPOs that feature the [BP2O83−] unit, may uncover interesting combinations of properties useful for the design of new functional open-framework materials. Herein, we report both the structure and properties of two new representatives that belong to the family of chiral MeBPOs of composition MIMII(H2O)2[BP2O8]·yH2O featuring MI = NH4+ and MII = Mn2+ and Co2+. The relationship of the new members within the family of related compounds is analyzed from both chemical and physical viewpoints. This systematic study intends to uncover key structure−property relationships in chiral BPOs with the ultimate goal to enable the targeted design of functional BPO materials.

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RESULTS AND DISCUSSION Crystal structure analysis reveals that I and II have (NH4)1−2xM1+x(H2O)2(BP2O8)·yH2O stoichiometry and belong to the MIMII(H2O)2[BP2O8]·yH2O family (Table 1). Table 1. Reported Representatives of the MIMII(H2O)2[BP2O8]·yH2O Family and Their Absolute Configurations

EXPERIMENTAL PART

Synthesis and Characterization. (NH4)1−2xM1+x(H2O)2(BP2O8)·yH2O, with M = Mn (I) and Co (II), were prepared by hydrothermal reactions. Typically, a mixture of M(OAc)2 ·4H2O (M = Co, Mn; 1 mmol, 0.25 g), H3BO3 (16.2 mmol, 1.0 g), NH4H2PO4 (3.5 mmol, 0.40 g), and demineralized water (55.6 mmol, 1.0 g) was loaded in a 15 mL Teflon-lined stainless steel autoclave and then heated under autogenous pressure at 200 °C for 7 days, followed by slow cooling to room temperature (RT) in 12 h. After washing with water and drying in air, almost colorless polycrystalline samples of I and rather large purple crystals of II (Figure S2) were obtained. Single-crystal X-ray diffraction (XRD) data were collected at RT and 100(2) K on a Bruker D8 Venture diffractometer equipped with a Mo Kα IμS microfocus source (λ = 0.71073 Å) and a Photon 100 CMOS detector. For data collection, the Bruker APEX3 program was used, and for data integration and reduction, the Bruker SAINT software package was utilized.61 The SADABS program was employed to treat the data for absorption effects by using the multiscan method.62 Initial models of the crystal structures were first obtained with the program SHELXT-201463 using direct methods and then refined using the program SHELXL-201464 within the APEX3 software package. Non-H atoms were refined anisotropically. Phase analyses were performed using powder XRD data collected at 293 K with the aid of a STOE STADI P powder diffractometer equipped with an area detector using Cu Kα1 radiation (λ = 1.54059 Å). The samples were dispersed on Mylar sheets, fixed with vacuum grease, and held in place with split aluminum rings. The lattice parameters were refined using the WinXPow program.65 Attenuated total reflectance (ATR) spectroscopy was carried out on an Agilent Technologies Cary 630 FT-IR spectrometer equipped with a diamond-crystal ATR unit. Solid samples were pressed on top of the ATR crystal with a diamond surface to ensure contact. Raman spectra were obtained at 150 mW on a Horiba Xplora Raman microscope at RT. Laser irradiation of 785 nm was used for excitation. Silicon was used as a standard for the calibration of Raman shifts.

MI/MII

Li

Na

H3O

K

NH4

Mn Fe Co Ni Cu Zn Mg Sc In

+

− − − −

−

− − ± −

+ + ±

− −

− −

+ + + − ± +

−

Ca

Ag

+ + −

+ +

− −

+

+ −

According to powder XRD patterns (Figure S3), both compounds were obtained as phase-pure. It is notable that the change in the unit cell parameters, most importantly in the c/a ratio, between compounds that contain H+, Li+, Na+, K+, or NH4+ is very small, while the influence of the transition metal is much more significant. For details on the data collection and structure analysis, see Table 2. A small, but notable change of the unit cell parameters and volume occurs at low temperatures. At 80 K, for I, a = 9.555(6) Å and c = 15.662(9) Å, while II exhibits a = 9.477(2) Å and c = 15.567(3) Å, compared to a = 9.6559(3) Å and c = 15.7939(6) Å for I at RT and a = 9.501(3) Å and c = 15.582(4) Å for II at RT. This corresponds to 3% and 0.5% volume reductions for I and II, respectively. The main structural features are helical chains of vertexsharing BO4 and PO4 tetrahedra featuring the BO4(PO4)2 fragment as the main building unit (Figures 1 and S1). The BO4 tetrahedra deviate only a little from the ideal with dB−O = ∼1.47 Å (Table 3), they are also connected exclusively to PO4 tetrahedra through vertex-sharing O atoms. The PO4 tetrahedra B

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N)1−x) with five shorter M−O contacts and a sixth contact to a mixed occupied site (Figure 2), which, in the case of an occupied M site, is the O atom of a water molecule or an ammonium cation when the M2 site is empty. Taking into account only the short M2−O contacts, the coordination polyhedron can best be described as a distorted tetragonal pyramid. Including the additional contact results in a severely distorted octahedral coordination by O atoms for M2. An ideal octahedral coordination in this case is impossible because of the steric restrictions evoked by the disordered water/ammonia position. The distorted M2 octahedron shares a trigonal face and a vertex with two neighboring M1O6 octahedra and may potentially share an edge with the identical M2O5 O polyhedron. The latter is, however, unlikely because of the low occupation of the M2 position. The M2 tetragonal pyramid shares an edge, vertex, and another vertex with the abovementioned neighbors, respectively. The ordered transitionmetal positions are well-separated (dM−M > 4 Å); however, the introduction of the extra disordered positions lowers this value to ∼3.1 Å, making magnetic ordering possible at lower temperatures (see below). The location and mutual orientation of the M2O6 and M1O6 octahedra, along with the low occupation of the M2 positions, may be the main reasons for H2O/NH4+ disorder. Total occupation of the M2 and N positions has been restricted to give a total charge of 1+ to balance the polyanionic network BP2O8 with an overall charge of 3−, while occupation of the M2 and disordered O positions has been refined independently, resulting in a slightly different composition in each case. Both Co2+- and Mn2+-containing compounds could be obtained in pure but different chiral forms, P6122 and P6522, respectively. In this light, it is interesting that two similar compounds with Co and Mn ions but H3O+ instead of NH4+ were reported with the P6 1 22 space group, 60 while (NH4)0.5Co1.25(H2O)2(BP2O8)·0.5H2O prepared from different precursors has been reported with the P65 space group; however, the correct determination of the space group is doubtful because perfect merohedral twinning [batch scaling factor (BASF) = 0.5] has been claimed.18 It is interesting that both B-containing precursors used for the reactions crystallize in enantiomorphic chiral space groups: B2O3 is known to crystallize in P31, while boric acid crystallizes in P32. This can provide some hint for the difference between two Co enantiomorphs obtained from B2O3 and H3BO3; however, two different results for Co and Mn ions with NH4+ and two identical results with H3O+ respectively point to a greater influence of both the cation and possibly reaction conditions. To date, almost 40 compounds are reported in the family, with nearly 50% in each chiral form (Table 1), while the Co/NH4 phase is a rare case, where both enantiomorphs have been reported. Thus, further investigations are required to understand this behavior. The IR spectra of I and II are shown in Figure S4. The bands at 3400 and 1620 cm−1 can be assigned to characteristic OH stretching and deformation vibrations, respectively. Bands centered at 1425 and 3100 cm−1 can be attributed to stretching and deformation vibrations of the NH groups. Bands in the range between 550 and 1200 cm−1 originate from PO4 and BO4 vibrations.66 Raman spectra of I and II in the range of 200− 2700 cm−1 (Figure S5) show quite similar features. The bands at 211 and 276 cm−1 can be attributed to MO stretching vibrations. The bands at 420 and 455 cm−1 are assigned to ν2 PO43− bending modes. The bands centered at 569 and 615

Table 2. Details of the Single-Crystal XRD Measurements and Data Collection of I and II empirical formula fw space group a, Å c, Å volume, Å3 Z density (calcd), g/cm3 μ, mm−1 F(000) Flack parameter θ range, deg index ranges reflns collected indep reflns refinement method data/restraints/ param GOF on F2 final R indices [I > 2σ(I)] R indices (all data) Rint largest diff peak and hole, e−/Å3

Mn1.25(NH4)0.50(BP2O8) (H2O)2.30 (I) 319.86 P6522 9.6559(3) 15.7939(6) 1275.3(1) 6 2.499

Co1.28(NH4)0.44(BP2O8) (H2O)2.41 (II) 327.68 P6122 9.501(3) 15.582(4) 1218.2(8) 6 2.680

2.329 952 0.01(3) 2.44−32.02 −14 ≤ h ≤ +13, −13 ≤ k ≤ +14, −23 ≤ l ≤ +23 18098 1444 full-matrix least

3.108 976 0.00(4) 4.64−30.81 −12 ≤ h ≤ +13, −13 ≤ k ≤ +10, −22 ≤ l ≤ +22 10317 1282 squares on F2

1444/0/79

1282/0/79

1.11 R1 = 0.029, wR2 = 0.080

1.07 R1 = 0.038, wR2 = 0.082

R1 = 0.032, wR2 = 0.082

R1 = 0.057, wR2 = 0.089

0.033 0.76 and −0.52

0.091 0.58 and −0.51

Figure 1. (top left) View along the [001] direction of anionic partial structures in I. (bottom left) View along the [001] direction showing the connection of MO6 octahedra in I and II. (right) View along the [010] direction of I and II illustrating the polyhedral connectivity.

are more distorted from the ideal symmetry, with the P−O distances ranging from 1.50 to 1.57 Å. However, it is the PO4 groups that connect the helical chains through metal cations. In the structures, there are two crystallographically independent metal positions, M1 and M2. For M1, an almost ideal octahedral coordination by O atoms is observed. The other M position (M2), which is only partially occupied by metal cations, features a 5 + 1 coordination according to (MO5(O/ C

DOI: 10.1021/acs.inorgchem.7b01423 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry Table 3. Atomic Positions and Equivalent Anisotropic Displacement Parameters of I and II x Mn1 Mn2 P1 B1 O1 O2 O3 O4 O5 O6 N6 Co1 Co2 P1 B1 O1 O2 O3 O4 O5 O6 N6

−0.45067(4) −0.9569(5) −0.83516(8) −0.1490(2) −0.9793(3) −0.8210(3) −0.6909(3) −0.8655(3) −0.5107(4) −0.9160(17) −0.855(2) 0.44894(5) 0.9559(7) 0.83191(14) 0.1510(4) 0.9784(4) 0.8203(4) 0.6846(4) 0.8618(4) 0.5178(6) 0.9096(16) 0.849(4)

y

z

Mn1.25(NH4)0.5(BP2O8)(H2O)2.30 −0.54933(4) −0.9167 −0.3222(5) −0.9084(3) −0.61936(8) −0.91869(4) −0.8510(2) −0.9167 −0.7902(2) −0.93393(12) −0.5918(2) −0.82062(11) −0.6207(3) −0.95407(13) −0.4935(3) −0.95550(13) −0.8094(4) −0.9523(2) −0.0840(17) −0.9167 0 0 Co1.28(NH4)0.44(BP2O8)(H2O)2.41 0.55106(5) 0.9167(5) 0.3247(7) 0.9102(3) 0.61567(15) 0.92046(7) 0.8490(4) 0.9167 0.7883(4) 0.93589(18) 0.5863(4) 0.82149(17) 0.6178(5) 0.95530(19) 0.4871(4) 0.95884(19) 0.8102(6) 0.9532(3) 0.0904(16) 0.9167 0 0

Ueq, Å2 0.0159(2) 0.0221(8) 0.0117(1) 0.0104(6) 0.0140(4) 0.0153(4) 0.0217(5) 0.0212(5) 0.0430(8) 0.095(5) 0.095(5) 0.0146(2) 0.026(1) 0.0124(3) 0.0117(14) 0.0150(7) 0.0137(7) 0.0196(8) 0.0189(8) 0.037(1) 0.099(7) 0.099(7)

SOF ≠ 1

0.125(1)

0.30(2) 0.5

0.14(1)

0.41(2) 0.44

Figure 2. Coordination polyhedra of the Mn positions in the crystal structure of I. Six-coordinated Mn sites are marked in blue and fivecoordinated in pink. O atoms are red and N atoms dark blue. Figure 3. RT UV−vis absorption spectrum of compound I.

cm−1 are likely to be due to the ν4 PO43− bending modes. The intense band located at 981 cm−1 is attributed to the ν1 PO43− symmetric stretching mode. The bands at 1048 and 1087 cm−1 correspond to the ν3 PO43− antisymmetric stretching mode.67 The high-resolution XPS Mn 2p spectrum of I displays two peaks at 642.4 and 654.7 eV for Mn 2p3/2 and Mn 2p1/2 (Figure S6), respectively, which are in agreement with Mn2+.68,69 The XPS Co 2p spectrum of II (Figure S6) shows two sharp peaks at binding energies of 781.9 eV (2p3/2) and 797.9 eV (2p1/2). Moreover, two satellites can be seen at 786.1 and 802.7 eV. These characteristic peaks correspond solely to Co2+.70,71 Thus, XPS data for both I and II provide conclusive evidence for the II+ oxidation states of Mn and Co in each compound. The optical properties of both compounds have been studied at RT and 77 K. The low-temperature absorption spectra of I and II are presented in Figures S7 and S8, respectively. The RT UV−vis absorption spectrum of I reveals several bands with maxima at 306, 341, 360, 406, 409, 423, and 538 nm (Figure 3). All of these bands are characteristic for Mn2+ ions in an octahedral crystal field.72 On the basis of the high-spin side of the d5 Tanabe−Sugano diagram, the seven absorption bands can be assigned to the following spin-forbidden transitions: 6 A1g(6S) → 4T1g(4P), 6A1g(6S) → 4Eg(4D), 6A1g(6S) →

T2g(4D), 6A1g(6S) → 4Eg(4G), 6A1g(6S) → 4A1g(4G), 6A1g(6S) → 4T2g(4G), and 6A1g(6S) → 4T1g(4G). The UV−vis−NIR absorption spectrum of II recorded at RT shows three bands at 512, 598, and 1069 nm (Figure 4). The second and third bands are split into two lines. The observed bands are expected for Co2+ ions in an octahedral crystal field.73 Using the high-spin side of the d7 Tanabe−Sugano diagram, the three absorption 4

Figure 4. RT UV−vis−NIR absorption spectrum of compound II ground with BaSO4 in a ratio of 1:1. D

DOI: 10.1021/acs.inorgchem.7b01423 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry bands can be ascribed to the following spin-allowed transitions: 4 T1g(4F) → 4A2g(4F), 4T1g(4F) → 4T1g(4P), and 4T1g(4F) → 4 T2g(4F). It is worth reminding that both the CoO6 and CoO5O octahedra are tetragonally elongated, however, to a different degree. Hence, the band splitting can be explained by lowering the symmetry from Oh to D4h. In D4h symmetry, the 4T2g first excited state splits into 4B2g and 4Eg levels, while the 4T1g second excited state splits into 4A2g and 4Eg ones. The band assignment has been supported by fitting the experimental band positions (in cm−1) to the energy levels on the appropriate Tanabe−Sugano diagram. This allowed for calculation of the crystal-field splitting energy (CFSE, Δ0) value and the Racah parameters (B and C). The obtained results are as follows: for I, Δ0 = 8155, B = 753 cm−1, and C = 3371 cm−1 (for a free Mn2+ ion, B = 860 cm−1 and C = 3850 cm−1); for II, Δ0 = 10200, B = 544 cm−1, and C = 2520 cm−1 (for a free Co2+ ion, B = 971 cm−1 and C = 4497 cm−1).74 The CFSE values imply that ligands surrounding Mn2+ and Co2+ ions produce a relatively weak field, so the d electrons of both metal ions adopt a highspin configuration. The Racah parameters for I and II are reduced to about 88 and 56% of the free ion values, respectively. This is a result of the nephelauxetic effect, which refers to the degree of orbital overlap in the metal−ligand σ bond. Generally, the stronger the reduction of the Racah parameters, the greater the nephelauxetic effect and the more covalent the character of the metal−ligand bonding. Thus, one can assume that Co−ligand bonds are more covalent than Mn− ligand ones, which is in agreement with the literature.75 These expectations are also confirmed by experimental data because the Co−ligand interatomic distances are shorter than the Mn− ligand ones (Table 4).

Figure 5. RT and low-temperature emission spectra of compound I excited at 409 nm.

coordinates were calculated (for orange, x = 0.54 and y = 0.36; for red, x = 0.66 and y = 0.30) and plotted on the CIE 1931 chromaticity diagram (Figure S10). The red shift of the emission band with decreasing temperature can be explained in the following way. As the temperature decreases, the unit cell volume also decreases (3% at 80 K), which results in an increase of the crystal-field strength acting on the Mn2+ ion. On the basis of the high-spin side of the d5 Tanabe−Sugano diagram, one can notice that, as the crystal-field strength increases the energy of the 4T1g(4G) → 6A1g(6S) transition decreases, which agrees with the measurements. The red shift with decreasing temperature has also been observed for other Mn compounds like manganese fluorinated selenite, Mn2(SeO3)F2.76 It is worth mentioning that the luminescence of the Mn2+ ion in Mn2(SeO3)F2 is practically quenched at 150 K, while in I, it remains up to RT. Luminescence decay curves of I (Figures S11 and S12) can be fitted by a monoexponential function, yielding lifetime values of 33.4 ± 0.3 μs at RT and 1.87 ± 0.01 ms at 77 K. An increase of the luminescence lifetime with decreasing temperature has also been reported for another manganese(II) borophosphate, KMnBP2O7(OH)2.31 Upon excitation at 578 nm at RT or 77 K, no luminescence of the octahedral Co2+ ion in compound II has been detected either in the visible region (610−800 nm) or in the NIR region (950−1700 nm). Magnetic susceptibilities of both samples follow the Curie− Weiss law in a wide range of measured temperatures, above 40 K for Co (II; Figure 6) and above 15 K for Mn (I; Figure 7). The Mn effective magnetic moment determined from the Curie−Weiss fit using the refined chemical formula is 5.92 μB/ Mn, in full quantitative agreement with the expected moment for Mn2+ (5.92 μB). The experimentally obtained Co moment, 5.1 μB/Co, is larger than the spin moment expected for the Co2+ ion (3.87 μB), yet it agrees very well with the magnetic moment values of open-framework phosphates,77 indicating the presence of an orbital moment. The Weiss temperatures are both negative, −8 K for I and −23 K for II, suggesting dominance of antiferromagnetic interactions in the ground state. At the same time, no clear signs of long-range magnetic ordering are seen down to 2 K for both compounds, although magnetization of I exhibits a sharp upturn and larger values compared to the same of II near 2 K.

Table 4. Interatomic Distances for I and II bond type

length (Å)

bond type

length (Å)

Mn1−O4 Mn1−O3 Mn1−O5 Mn2−O5 Mn2−O6 Mn2−N6 Mn2−O5 Mn2−O4 Mn2−O3 Mn2−O5 Mn1−Mn2 P1−O4 P1−O3 P1−O1 P1−O2 B1−O1 B1−O2

2.124(2) 2.146(2) 2.346(3) 2.059(6) 2.13(2) 2.28(2) 2.348(5) 2.351(5) 2.370(5) 3.068(5) 3.182(4) 1.503(2) 1.507(2) 1.556(2) 1.566(2) 1.464(3) 1.467(3)

Co1−O4 Co1−O3 Co1−O5 Co2−O5 Co2−O6 Co2−N6 Co2−O5 Co2−O4 Co2−O3 Co2−O5 Co1−Co2 P1−O4 P1−O3 P1−O1 P1−O2 B1−O1 B1−O2

2.034(3) 2.087(4) 2.281(4) 2.017(8) 2.04(2) 2.25(2) 2.303(7) 2.266(6) 2.285(6) 2.952(7) 3.103(6) 1.509(4) 1.511(4) 1.550(4) 1.561(3) 1.472(3) 1.467(5)

The luminescent properties of compound I have also been investigated. The excitation spectra of I at RT and 77 K (Figure S9) are consistent with the corresponding UV−vis absorption spectra. Upon excitation at 409 nm, compound I shows a broad emission band centered around 635 nm at RT, which corresponds to the spin-forbidden 4T1g(4G) → 6A1g(6S) transition of the Mn2+ ion (Figure 5). At 77 K, the band is shifted by 20 nm and its maximum is around 655 nm. As a result, the emission color changes from orange at RT to red at 77 K. To visualize the emission color change, the chromaticity E

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Inorganic Chemistry

Figure 6. Temperature dependence of the molar magnetic susceptibility, χmol, and inverse molar magnetic susceptibility (inset) for Co1.28(NH4)0.44(BP2O8)(H2O)2.41 measured in a 0.1 T applied magnetic field.

Figure 8. Magnetic-field dependence of isothermal magnetization measured for Co- and Mn-based BPOs at selected temperatures.

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CONCLUSIONS

Manganese and cobalt borophosphates of the composition (NH4)1−2xM1+x(H2O)2(BP2O8)·yH2O have been obtained using hydrothermal methods, and their structures have been determined by means of single-crystal XRD. Although both compounds are nearly identical in many aspects, they crystallize in enantiomorphic space groups P6522 and P6122 (a = 9.50− 9.66 Å, c = 15.8−15.6 Å, and V = 1218−1275 Å3). Manganese borophosphate shows a noticeable thermal expansion between 80 and 300 K that is 6 times larger compared to the analogous Co compound. Both compounds feature mixed helical anionic networks of PO4 and BO4 groups sharing common vertices and connecting via terminal O atoms to the transition metals in two different coordination environments and disordered chains of water and ammonia. The absorption spectra demonstrate characteristic bands for octahedral Mn2+ and Co2+ ions. The CFSE values suggest that both metal ions are situated in the weak crystal field. The Racah parameters for both compounds are reduced because of the nephelauxetic effect. Because the reduction is more pronounced for II than for I, the nature of the Co−ligand bonding is more covalent than of the Mn−ligand one, which also is reflected by its higher expansion coefficient. Furthermore, compound I reveals interesting luminescent properties. In contrast to many Mn compounds, of which the luminescence is compromised because of the concentration and temperature quenching effects, I exhibits an orange emission at RT. The emission color of I can be changed from orange to red by lowering the temperature from RT to 77 K. As the temperature decreases, the luminescence lifetime of I becomes longer. It increases from 33.4 μs at RT to 1.87 ms at 77 K. Both compounds are Curie− Weiss paramagnets with effective magnetic moments corresponding to those of noninteracting Mn2+ (Co2+) ions and a sizable orbital moment contribution in II. Negative Weiss temperatures are indicative of prevailing antiferromagnetic interactions in the ground states, although no long-range magnetic ordering has been detected down to 2 K.

Figure 7. Temperature dependence of the molar magnetic susceptibility, χmol, and inverse molar magnetic susceptibility (inset) for Mn1.25(NH4)0.5(BP2O8)(H2O)2.3 measured in a 0.1 T applied magnetic field.

Negative deviations from the Curie−Weiss behavior observed at low temperatures for both I and II indicate enhanced low-temperature magnetization and likely short-range magnetic correlations. The magnetic field dependence of magnetization measured at 2 K (Figure 8) confirms that both materials are no longer truly paramagnetic at this temperature. In fact, the Mn compound shows a tendency toward saturation above ∼2 T even though the magnetization increases slowly. The value of the magnetic moment (μs = 3.2 μB/Mn) is significantly lower than the theoretical net spin moment expected for Mn2+ (5 μB) possibly because of a lack of longrange magnetic order (the magnetization keeps increasing with B). The Co compound shows similar behavior but with weaker saturation tendency and lower magnetic moment. It is possible that both compounds may order magnetically below 2 K. However, we do not have sufficient evidence to assert whether their ground states are conventional long-range-ordered or are those of magnetically frustrated systems, like superparamagnetism, which has been reported for some open-framework metal phosphates.77 F

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.7b01423. Crystal photograph of II, powder XRD patterns, XPS, luminescence, IR, and Raman spectra for both compounds (PDF) Accession Codes

CCDC 1551772−1551773 contain the supplementary crystallographic data for this paper. These data can be obtained free of charge via www.ccdc.cam.ac.uk/data_request/cif, or by emailing [email protected], or by contacting The Cambridge Crystallographic Data Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax: +44 1223 336033.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Volodymyr Smetana: 0000-0003-0763-1457 Yaroslav Mudryk: 0000-0003-2658-0413 Anja-Verena Mudring: 0000-0002-2800-1684 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This research was supported by Grant VR 2016-05404 (to A.V.M.). Magnetic measurements (by Y.M. and V.K.P.) were supported by the U.S. Department of Energy, Office of Basic Energy Sciences, Division of Materials Sciences and Engineering (measurements of the magnetic properties). Ames Laboratory is operated for the U.S. Department of Energy by Iowa State University under Contract DE-AC02-07CH11358.

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