Ammine-Stabilized Transition-Metal Borohydrides of Iron, Cobalt, and

Oct 21, 2015 - Complex metal hydrides have been intensively studied as materials for energy storage because of their high gravimetric and volumetric h...
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Ammine-Stabilized Transition-Metal Borohydrides of Iron, Cobalt, and Chromium: Synthesis and Characterization Elsa Roedern and Torben R. Jensen* Interdisciplinary Nanoscience Center (iNANO) and Department of Chemistry, Aarhus University, Langelandsgade 140, DK-8000 Aarhus C, Denmark S Supporting Information *

ABSTRACT: Iron and cobalt borohydrides stabilized by ammonia (NH3), [Fe(NH3)6](BH4)2 and [Co(NH3)6](BH4)2, were synthesized along with a solid solution, [Co(NH3)6](BH4)2−xClx (x ∼ 1), and a bimetallic compound, [Fe(NH3)6](Li2(BH4)4). The compounds were prepared by new lowtemperature, solvent-based synthesis methods, using dimethyl sulfide or liquid NH3, which allow for the removal of inert metal halides. The crystal structures were determined from synchrotron radiation powder X-ray diffraction (SR-PXD) data. [M(NH 3 ) 6 ](BH 4 ) 2 (M = Fe, Co) and [Co(NH 3 ) 6 ](BH4)2−xClx crystallize in the cubic crystal system, where the transition metals are octahedrally coordinated by NH3. Polymeric chains of lithium coordinated by four bridging BH4− anions are found in [Fe(NH3)6](Li2(BH4)4). The new compounds have high hydrogen densities of ∼14 wt % H2 and ∼140 g H2/L and release a mixture of hydrogen and NH3 gas at low temperatures, T < 80 °C. The decomposition mechanisms of the prepared compounds along with the composites [Fe(NH3)6](BH4)2·nNH3BH3 (n = 2, 4, 6) were studied by thermal analysis and in situ SR-PXD.

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INTRODUCTION Complex metal hydrides have been intensively studied as materials for energy storage because of their high gravimetric and volumetric hydrogen content.1−4 A drawback for many materials are the unfavorable thermodynamic and kinetic properties, leading to high hydrogen release temperatures unsuitable for proton-exchange membrane fuel-cell applications.5,6 The thermodynamic properties of too stable metal borohydrides may be improved by the formation of reactive hydride composites,4,7,8 while nanoconfinement or eutectic melting of metal borohydrides mainly improves the kinetic properties.9−12 An empirical correlation between the stability of metal borohydrides and the Pauling electronegativity, χP, of the metal coordinated by BH4− complexes is well established.5,13 High electronegativity leads to low decomposition temperatures, and metals with χP > ∼1.5 are often reduced and release both hydrogen and diborane.14 However, transition-metal borohydrides with d-electron configurations d0, d5, or d10 are stable at room temperature (RT) and may be obtained as halide-free bimetallic borohydrides by other new synthesis methods.4,15 Metal reduction can successfully be prevented by introducing neutral ligands like ammonia (NH3), demonstrated by the synthesis of ammine transition-metal borohydrides of titanium (d1) and vanadium (d2).16,17 Compounds based on the most abundant transition metals, chromium, iron, cobalt, nickel, and copper, have not been characterized in any detail since the initial work in the 1960s and 1970s18,19 because of © XXXX American Chemical Society

their low thermal stability. We present a new solvent-based lowtemperature method to synthesize transition-metal borohydrides stabilized by NH3 to overcome the challenge of their low stability and focused our investigations on the transition metals chromium, iron, cobalt, nickel, copper, and zinc.

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EXPERIMENTAL DETAILS In the first method (Scheme 1), a salt metathesis reaction between a transition-metal chloride and an alkali-metal borohydride in precooled anhydrous dimethyl sulfide, S(CH3)2, produces a solution containing the transition-metal borohydride, which can be separated from the solid byproducts and remaining reactants and then reacted with NH3 to produce ammine transition-metal borohydride. MCl2 (M = Cr, Fe, Co, Ni; anhydrous, 99.99%) and LiBH4 (95%) were reacted in the molar ratio 1:2 in precooled (T ∼ −50 °C) anhydrous S(CH3)2. The reaction mixture was stirred for 4 h and then filtered, to remove the produced LiCl and unreacted MCl2 and LiBH4. The filtrate was immediately reacted with NH3 gas, and a precipitate formed. After 20 min, the excess NH3 and the solvent were removed under dynamic vacuum. During the entire synthesis, the flask was kept in a ethanol−dry ice cooling bath, at T < −30 °C. The samples were stored in a glovebox freezer at −30 °C and are denoted M_1. Received: August 25, 2015

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DOI: 10.1021/acs.inorgchem.5b01959 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry Scheme 1. Synthesis in Organic Solvent

Scheme 2. Synthesis in Liquid NH3

Hiden Analytical HPR-20 QMS sampling system. Samples of approximately 3 mg were placed in an aluminum crucible and heated from 15 to 500 °C at a heating rate of 5 °C/min in an argon flow (20 mL/min). The evolved gas was transported to the mass spectrometer through heated tubing. In Situ Time-Resolved SR-PXD (In Situ SR-PXD). For PXD analysis, the samples were packed into precooled 0.5 mm borosilicate or quartz capillaries. All tools were stored and cooled to −30 °C in a glovebox freezer and taken out for no more than 1−2 min at a time. The capillaries containing the samples were stored in the freezer. The reaction pathway of the decomposition was investigated using in situ SR-PXD. Highresolution SR-PXD data of samples Cr_2, Fe_2, and Co_2 were obtained at −33 °C at Beamline I11 at the Diamond Synchrotron, Didcot, U.K., using a multianalyzing crystal (MAC) detector with five arms with 45 MAC channels each and a wavelength of 0.8258 Å. In situ SR-PXD data were collected at the same beamline for Cr_1 and Fe_1, using a PSD detector and a cryostream for heating, and at Beamline I711 at the MAX IV laboratories in Lund, Sweden, for samples Fe_2 and Co_2 with a MAR165 CCD detector system with an X-ray exposure time of 30 s and a selected wavelength of λ = 0.9938 Å. Structure Solution. The structures of all new compounds were refined using the Rietveld method, implemented in the program Fullprof.20 The final refinements are shown in the Supporting Information, Figures S6−S9. The individual isotropic B factors were not refined but fixed to the following values: Biso(M = Fe, Co) = 1.0, Biso(Cl) = 2.0, Biso(Li) = 3.0, Biso(B) = 3.0, Biso(N) = 3.0, and Biso(H) = 5.0. Boverall was refined for all structures.

For cobalt and nickel, a black coloring of the solution was observed during the reaction time (4 h), and only ammonia borane (NH3BH3) was produced, identified by IR and powder X-ray diffraction (PXD). At RT, LiBH4 dissolved in S(CH3)2 reduces MCl2 (M = Cr, Fe, Co, Ni, Zn) and CuCl, observed as a black or characteristic metallic coloring of the solution or glassware from metal M, and a S(CH3)2 BH3 complex solution is produced during the reaction. The reaction with NH3 gas yields NH3BH3 and not a metal borohydride. Thus, the solvent S(CH3)2 successfully prevents the release of highly toxic and flammable diborane gas, B2H6, making the synthesis method safer. At low temperatures, the reduction of the metal is hindered, and M(BH4)2−S(CH3)2 solutions could be obtained for M = Cr or Fe. In the second method (Scheme 2), the metathesis reaction between a transition-metal chloride and an alkali-metal borohydride is carried out in liquid NH3 to directly produce an ammine transition-metal borohydride. Solid metal chloride MCl2 (M = Cr, Fe, Co) or CuCl (anhydrous, 99.99%) and LiBH4 (95%) were loaded in the molar ratio 1:2 (1:1 for Cu) into a flask and cooled in an ethanol−dry ice bath to T ∼ −50 °C. NH3 gas was condensed into the flask at −50 °C, and the reaction mixture was stirred for 4 h. The flask was then allowed to warm to −25 °C, causing NH3 to boil off. Excess gas was removed under vacuum, and the samples were stored in a glovebox freezer at −30 °C and are denoted as M_2 (M = Cr, Fe, Co). For the copper sample, synchrotron radiation powder X-ray diffraction (SR-PXD) revealed the presence of metallic α-Cu and weak, unidentified diffraction peaks. All handling and manipulation of the chemicals were carried out under an argon atmosphere in an MBraun Unilab glovebox with a recirculation gas purification system and gas/humidity sensors [p(O2,H2O) < 1 ppm] or by using Schlenk techniques. All commercial chemicals (Sigma-Aldrich) were used as received. The obtained products were characterized by SRPXD, thermal analysis (TGA−DSC−MS), and IR spectroscopy. Fourier Transform Infrared Spectroscopy (FTIR). FTIR spectra were measured using a Nicolet 380 Avatar spectrometer with a diamond ATR accessory. For the IR measurements, a sample was put into a precooled glass vial kept at −30 °C until right before the measurement. The sample was transferred onto the diamond, pressed, and immediately measured. Short exposure to air during this procedure, as well as warming of the sample during the measurement, cannot be avoided; however, the spectra did not change during the data collection time of 30 s. TGA−DSC−MS. The decomposition reactions were studied by combined thermogravimetric analysis (TGA), differential scanning calorimetry (DSC), and mass spectrometry (MS) of the evolved gas, using a PerkinElmer STA 6000 apparatus and a

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RESULTS AND DISCUSSION Synthesis. The reaction of FeCl2 with LiBH4 in S(CH3)2 at low temperatures (Scheme 1, Fe_1) provides a clear solution of Fe(BH4)2, which was reacted with NH3 gas and dried under vacuum to produce a white solid of [Fe(NH3)6](BH4)2. [Fe(NH3)6](BH4)2 was also successfully synthesized in liquid NH3 (Scheme 2, Fe_2). The new compound [Co(NH3)6](BH4)2 was synthesized in liquid NH3, Co_2, but was reduced to cobalt in S(CH3)2 (Co_1). The reaction between CrCl2 and LiBH4 in S(CH3)2 at low temperatures produces a green solution, and a purple precipitate upon reaction with NH3 (Cr_1). An overview of the composition of the crystalline products is provided in Table 1. Structures. A total of five new compounds are identified, structurally characterized, and reported in Table 2. The new cubic compounds [Co(NH3)6](BH4)2 and [Fe(NH3)6](BH4)2 are built from octahedral complexes of [M(NH3)6]2+ with BH4− acting as counterions (see Figure 1) and are isostructural to the halides [Co(NH3)6]Cl2 and [Fe(NH3)6]Cl2. The unit cell of [Fe(NH3)6](BH4)2 (Fe_1 and Fe_2), a = 10.7142(1) Å, B

DOI: 10.1021/acs.inorgchem.5b01959 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry Table 1. Synthetic Scheme, Reactants, and Composition of the Crystalline Products Extracted by Rietveld Refinement of SR-PXD Dataa sample

reactants

products

Fe_1

FeCl2 + LiBH4 in S(CH3)2 + excess NH3 FeCl2 + LiBH4 in liquid NH3 CoCl2 + LiBH4 in S(CH3)2 + excess NH3 CoCl2 + LiBH4 in liquid NH3

[Fe(NH3)6](BH4)2, 73 wt %; [Fe(NH3)6] (Li2(BH4)4), 11 wt %; Li(NH3)BH4, 16 wt %

Fe_2 Co_1 Co_2 Cr_1 Cr_2

CrCl2 + LiBH4 in S(CH3)2 + excess NH3 CrCl2 + LiBH4 in liquid NH3

[Fe(NH3)6](BH4)2, 59 wt %; [Fe(NH3)6]Cl2, 16 wt %; LiCl, 25 wt % NH3BH3b [Co(NH3)6](BH4)2, 40 wt %; [Co(NH3)6]Cl2, 21 wt %; [Co(NH3)6](BH4)2−xClx, 16 wt %; LiCl, 8 wt %; LiCl·NH3, 15 wt % u1, u2 u3

a

u1, u2, and u3 denote unidentified products. PXD data and IR spectra are shown in Figures S1 and S2. bIdentified by in-house PXD and IR spectroscopy.

Figure 1. Crystal structure of cubic [M(NH3)6](BH4)2, where M = Co, Fe. [M(NH3)6]2+ are blue octahedra, BH4− anions are red tetrahedra, and hydrogen atoms are not shown.

is slightly larger than that of [Fe(NH3)6]Cl2, a = 10.1899 Å, and the Fe−N bond distances are 2.2379 and 2.2065 Å, respectively. [Co(NH3)6](BH4)2 (Co_2) also has a cubic structure, a = 10.6724(3) Å, and a Co−N bond distance of 2.1836 Å and is isostructural to [Fe(NH3)6](BH4)2 and [Co(NH3)6]Cl2, a = 10.1774 Å.21 A new tetragonal bimetallic ammine iron lithium borohydride, [Fe(NH3)6](Li2(BH4)4), was identified as a minor crystalline component (10 wt %) in Fe_1, which is isostructural to [Mn(NH3)6](Li2(BH4)4).22 The new compound is built from [Fe(NH3)6]2+ octahedra and polymeric chains along [001] of lithium coordinated by four bridging BH4− anions, formally forming complex anions [Li2(BH4)4]2− (see Figure 2). The formation of [Fe(NH3)6](Li2(BH4)4) may be caused by incomplete salt metathesis in the first step of reaction 1 and the carryover of some LiBH4. Another set of weak diffraction peaks in this sample is assigned to ammine lithium borohydride, Li(NH3)BH4. Furthermore, a new solid solution, [Co(NH3)6](BH4)2−xClx, formed from unreacted CoCl2 and the ammine cobalt borohydride is observed in sample Co_2. The structural model of [Co(NH3)6](BH4)2 was used as a starting model for the solid solution [Co(NH3)6](BH4)2−xClx. Chloride substitution was allowed at the boron site in the [Co(NH3)6](BH4)2

Figure 2. Crystal structure of [Fe(NH3)6](Li2(BH4)4). [Fe(NH3)6]2+ shown as blue octahedra and lithium (green) coordinated by four bridging BH4− complexes (red tetrahedra), forming a polymeric chain of [Li2(BH4)4]2−.

structure, and the occupancy was refined to Occ(Cl) = x = 0.96(2), using the constraint Occ(B) + Occ(Cl) = 2. The unit cell parameter of the solid solution, a = 10.4417(9) Å, is

Table 2. Structural Data and Hydrogen Content for the New Compounds crystal system space group a (Å) c (Å) Z V (Å3) M (g/mol) ρ(H2) (wt %) ρ(H2) (kg of H2/m3) χP (M) T (°C)c a

[Fe(NH3)6](BH4)2a

[Fe(NH3)6](Li2(BH4)4)

[Co(NH3)6](BH4)2

[Co(NH3)6](BH4)2−xClxb

LiCl·NH3

u1

cubic Fm3̅m 10.7142(1)

tetragonal P42/mnm 10.7117(4) 6.8963(4) 2 791.29(6) 231.3 14.8 143.8 1.83 (Fe)/ 0.98 (Li) −33

cubic Fm3̅m 10.6724(3)

cubic Fm3̅m 10.4417(9)

cubic Pm3̅m 3.8495(1)

hexagonal

4 1215.57(6) 190.8 13.7 143.1 1.88 −33

4 1138.43(14) 211.4 10.5 130.0 1.88 −33

1 57.04(1) 59.425 5.1 76.2 0.98 −33

4 1229.93(1) 187.7 14.0 141.6 1.83 −33

6.559(1) 6.179(1) 230.21(1)

1.66 37

Results of refinement of Fe_1. bx = 0.96 at −33 °C. cTemperature of data collection. C

DOI: 10.1021/acs.inorgchem.5b01959 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry between the values of the ammine metal chloride and metal borohydride, thus obeying Vegard’s law. A set of Bragg peaks observed in Co_2 was assigned to another new cubic compound, LiCl·NH3 (a = 3.8495(1) Å) isostructural and isoelectronic to LiCl·H2O (a = 3.83 Å).23 A more detailed structural analysis of LiCl·NH3 may reveal a lower symmetry, similar to LiCl·H2O, where tetragonal and orthorhombic superstructures have been described.24 The samples Cr_1 and Cr_2 contain several crystalline unidentified compounds (denoted as u1, u2, etc.), and decomposition-aided indexing of in situ SR-PXD data of Cr_1 provide a hexagonal unit cell of u1. Thermal Decomposition. The stability and decomposition mechanisms of the new compounds were investigated using in situ SR-PXD and TGA−DSC−MS. In situ PXD data of Fe(BH4)2(NH3)6 (Fe_1; see Figure 3) reveal decomposition at

Figure 4. TGA−DSC−MS of samples Cr_2, Fe_2, and Co_2. ΔT/Δt = 5 °C/min in argon flow. Top: TGA data. Bottom: MS data of Cr_2, Fe_2, and Co_2.

(see Figure S5). The sample Cr_1 releases gas continuously in the temperature range of RT to 500 °C for a total mass loss of 38 wt % (Figure S5). MS data reveal that the release of diborane, B2H6, is not observed. As discussed above, the transition-metal borohydrides M(BH4)2 (M = Cr, Fe, Co, Ni) are not stable at RT and LiBH4 dissolved in (CH3)2S readily reduces MCl2 (M = Cr, Fe, Co, Ni) and CuCl, observed as black or metallic coloring of the solution. The transition metals in [M(NH3)6](BH4)2, M = Co or Fe, and [Fe(NH3)6](Li2(BH4)4) are protected against reduction by coordination to NH3, which stabilizes the solids at RT. Their decomposition appears to be initiated by NH3 release, which may allow metal cation reduction, possibly accompanied by the release of hydrogen and B2H6. The emission of B2H6 may not be observed because of immediate reaction with the remaining NH3. This mechanism is supported by the observation of metallic iron and cobalt by in situ SRPXD at higher temperatures, while the boron- and nitrogencontaining compounds are X-ray amorphous. Thus, this work suggests that diborane release may be suppressed by NH3 in amine metal borohydrides, similar to the addition of metal hydrides in manganese borohydride.26 Furthermore, we prepared the ammine iron borohydride− ammonia borane composites Fe(NH3)6(BH4)2·nNH3BH3, where n = 2, 4, 6, and analyzed the gas release. The two components Fe(NH3 ) 6 (BH 4 ) 2 and NH 3 BH 3 appear to decompose individually, and the hydrogen content was not improved. This is further discussed in the Supporting Information and Figure S10. NH3 release from octaammine zirconium borohydride appears to be somewhat suppressed in composites containing ammonia borane, Zr(NH3)8(BH4)4·

Figure 3. In situ SR-PXD of [Fe(NH3)6](BH4)2 (Fe_1) from −33 to 77 °C. ΔT/Δt = 5 °C/min, p(Ar) = 1 bar, and λ = 0.825645 Å. Symbols: 1, [Fe(NH3)6](BH4)2; 2, Li(NH3)BH4; 3, [Fe(NH3)6](Li2(BH4)4).

∼63 °C, while diffraction from [Fe(NH3)6](Li2(BH4)4) and Li(NH3)BH4 disappears simultaneously at ∼40 °C, below the reported melting temperature of Li(NH3)BH4 at 57 °C.25 In situ PXD data of Co_2 (Figure S3) reveal four crystalline compounds at −33 °C: [Co(NH3)6](BH4)2, [Co(NH3)6]Cl2, LiCl·NH3, and LiCl. The weak diffraction of LiCl·NH3 disappears already at ca. −15 °C, while [Co(NH3)6](BH4)2 decomposes at ∼44 °C and [Co(NH3)6]Cl2 at 150 °C via the intermediate [Co(NH3)2]Cl2 observed in the temperature range 150−230 °C. Cobalt crystallizes at 150 °C, and the diffracted intensity of LiCl increases. A high background at low Bragg angles in the in situ SR-PXD data (Figure S4) suggests partial decomposition of sample Cr_1. Two unidentified compounds, denoted as u1 and u2, are observed at −33 °C. u1 was indexed in a hexagonal unit cell with a = 6.559(1) Å and c = 6.179(1) Å and decomposes at 27 °C, while the diffracted intensity of u2 more slowly decreases in the temperature range of ∼27 to ∼50 °C. TGA−DSC−MS of the released gas for samples Cr_2, Fe_2, and Co_2 (Figure 4) reveals total mass losses of 34, 35, and 29 wt %, respectively, in the temperature range of RT to 300 °C. The halide-free [Fe(NH3)6](BH4)2 (Fe_1) releases a mixture of hydrogen and NH3 (35 wt %) in a single step below 100 °C D

DOI: 10.1021/acs.inorgchem.5b01959 Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry

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nNH3BH3 (n = 2, 3, 4, 5),27 but release of toxic gases, e.g., B2H6 and B3N3H6, from M(BH4)n−NH3BH3 (M = Li, Na, K, Mg, Ca) is not suppressed.28−32 Even though chromium has the lowest electronegativity of chromium, iron, and cobalt, the respective chromium complex appears to be less stable. This may be due to the less favorable octahedral coordination of chromium(II) to NH3 similar to [Cr(NH3)6]X2 (X = Cl, Br, I), which readily lose NH3 and form pentaammines, while [M(NH3)6]X2 (M = Co, Fe; X = Cl, Br, I) compounds are more stable.21,33

CONCLUSION In summary, we have described a new synthesis strategy developed to target instable transition-metal borohydrides, e.g., iron the second most abundant metal on earth, which is of particular interest for energy applications. Using organic solvents allows for the removal of inert metal halides, and at low temperatures, mono- or bimetallic borohydrides stabilized by NH3 can be synthesized, which are not accessible by mechanochemistry and other solvent-based methods.15,34,35 [Fe(NH3)6](BH4)2, [Co(NH3)6](BH4)2, the solid solution [Co(NH3)6](BH4)2−xClx, and bimetallic compound [Fe(NH3)6](Li2(BH4)4) were structurally and physically characterized. The new materials have very high hydrogen densities and low decomposition temperatures, onset below 80 °C; however, further work is needed to suppress NH3 release. The new synthetic routes presented here may provide catalytic transition metal and reducing borohydride solutions valuable as reagents for organic transformations or precursors for novel functional materials, which may combine hydrogen storage, ion conduction, and optical or magnetic properties.36−39 ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.5b01959. CIF files for all new crystal structures (ZIP) Supplementary figures (PDF)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Author Contributions

The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the Innovation Fund Denmark (Project HyFill-Fast), the Danish National Research Foundation, Center for Materials Crystallography (DNRF93), and the Danish Research Council for Nature and Universe (Danscatt). Access to beamtime at Beamline I11 at the Diamond Synchrotron, Didcot, U.K., and at Beamline I711 at the MAX-II Synchrotron, Lund, Sweden, in the research laboratory MAX-lab is gratefully acknowledged. We are grateful to the Carlsberg Foundation. E

DOI: 10.1021/acs.inorgchem.5b01959 Inorg. Chem. XXXX, XXX, XXX−XXX

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DOI: 10.1021/acs.inorgchem.5b01959 Inorg. Chem. XXXX, XXX, XXX−XXX