msceJournal of Materials Science and Chemical Engineering2327-60532327-6045Scientific Research Publishing10.4236/msce.2026.1410002msce-154435ArticleChemistryMaterials ScienceHigh-Purity Amorphous Boron Synthesis from Boron Trifluoride and Lithium Hydride0000-0001-6546-8526NikoleishviliPaata1TsurtsumiaGigla1LezhavaTina1KoiavaNana1KurtanidzeRusudan1JokhadzeNunu1 R. Agladze Institute Inorganic Chemistry and Electrochemistry, Ivane Javakhishvili Tbilisi State University, Tbilisi, Georgia
The authors declare no conflicts of interest regarding the publication of this paper.
Ultrapure amorphous elemental boron was prepared via a two-stage route starting from boron trifluoride diethyl etherate (BF3·Et2O) and lithium hydride. In the first stage, LiBH4 was synthesized in tetrahydrofuran at 50˚C - 55˚C using an optimized LiH:BF3 molar ratio of 4.7:1, yielding 93% isolated product at >99% 11B NMR purity, fluoride-free and with negligible diborane emission. In the second stage, the isolated LiBH4 was thermally decomposed in a horizontal tube furnace under argon at 650˚C, with thermal events consistent with decomposition via a Li2B12H12 intermediate. Staged ethanolysis and dilute acid purification yielded amorphous nanostructured boron of 98.7% purity, with primary particle size ~60 nm, specific surface area 45 m2·g−1.
Ultrapure amorphous elemental boron is a material of considerable technological importance for aerospace engineering, nuclear technologies, and high-energy-density fuels, but its direct synthesis from the elements is impractical. The direct reaction of lithium, boron, and hydrogen to form LiBH4—a necessary intermediate for elemental boron production—requires extreme conditions (150 bar H2, 650˚C), and the subsequent thermal decomposition of this intermediate to release elemental boron is itself a multi-step process requiring temperatures above 600˚C. Moreover, elemental boron is not cheaply available, so indirect routes via lithium borohydride are used instead [1][2]. The wet-chemical route through the reaction of lithium hydride with boron trifluoride in ethereal solvents has therefore become the dominant industrial method for the synthesis of lithium borohydride, with the LiBH4 subsequently decomposed thermally to release amorphous boron.
Lithium borohydride was first synthesized in 1940 by Schlesinger and Brown [1] by the reaction of ethyllithium with diborane but attempts to prepare it via the alternative reaction of LiH with BF3 in ethereal solvents began somewhat later. Elliott, Boldebuck and Roedel [3] first reported the reaction of LiH with boron trihalide ether complexes—including BF3·Et2O—to generate diborane, using LiH:BF3 molar ratios between 2:1 and 3.7:1; their experimental focus was diborane yield and purity, and no concrete data on LiBH4 formation or yield were given. Fedneva [4] subsequently investigated the mechanism of this reaction in detail and concluded that LiBH4 is not formed directly and stably in a single step, contrary to Elliott’s assumption. Instead, diborane forms first as the primary boron-containing product
4LiH + BF3·Et2O → LiBH4 + 3LiF↓ + Et2O↑ (1)
and is only subsequently able to react further with excess LiH in a second step to build up lithium borohydride. A follow-up communication by Fedneva [4] examined the reaction at a 4:1 LiH:BF3 molar ratio near 0˚C and found that, owing to the comparatively high solubility of LiF under these conditions, a large amount of lithium fluoroborate (LiBF4) forms as a competing byproduct while very little LiBH4 is actually produced—demonstrating that stoichiometry alone is insufficient for good yield, and that reaction temperature, LiF solubility, and reagent addition order are all critical variables [4].
These mechanistic constraints were resolved industrially by Hauk and Wietelmann [5] (Chemetall GmbH), who patented a process specifying a LiH:BF3 molar ratio greater than 4.1:1 (particularly preferably >4.4:1), finely divided LiH (mean particle size d50 ≤ 0.2 - 0.5 mm), a small LiBH4 seed charge (0.5 - 10 mol%) to reliably initiate the exothermic reaction, and reaction in a higher-boiling ethereal solvent—tetrahydrofuran, 2-methyltetrahydrofuran, tetrahydropyran, or 1,2-dimethoxyethane—at 10˚C - 70˚C, preferably 30˚C - 70˚C, in “semi-batch” mode with LiH suspended first and BF3 metered in [2][6]. Under these conditions the patent reports LiBH4 yields ≥90% (based on BF3) with purity >95% by 11B-NMR, essentially free of fluoride and with negligible diborane emission; worked examples achieved 93% and ~95% isolated yield in THF and dimethoxyethane, respectively, versus only 46% yield and heavy LiB2H7/BH3 contamination when the LiH:BF3 ratio fell to 4.08:1 [5]. More broadly, alongside mechanochemical ball-milling routes from borate precursors, the reaction of BF3 with LiH in ethereal solvents remains one of the principal industrial methods by which LiBH4 is produced today.
Once synthesized, LiBH4 can be thermally decomposed to release elemental boron. LiBH4 is a salt-like, hygroscopic crystalline material adopting the orthorhombic Pnma space group at room temperature (a = 7.173 Å, b = 4.434 Å, c = 6.798 Å, ρ = 0.669 g·cm−3) and undergoing a polymorphic transformation to a disordered high-temperature hexagonal phase near 108˚C, melting near 268˚C - 270˚C [1][7]. The overall decomposition reaction
LiBH4 → LiH + B + (3/2) H2 (2)
was long recognized to proceed through several intermediate steps rather than in one clean transition, with 13.8 wt% hydrogen released only when heated to 900˚C [1]. Orimo et al. [8] provided early spectroscopic evidence for this complexity: room-temperature Raman measurements on LiBH4 heated above 430˚C revealed spectral signatures corresponding to both an intermediate Li2B12H12 phase and amorphous boron as coexisting decomposition products, and combined with first-principles phase-stability calculations this work established Li2B12H12 as the thermodynamically most stable intermediate compound on the pathway to elemental boron, proposing two competing reactions with distinct mass losses (≈10 wt%, 56 kJ/mol H2 via the Li2B12H12 route; a further ≈4 wt%, 125 kJ/mol H2 toward LiH + B). Hwang et al. [9] gave direct 11B MAS/CPMAS NMR confirmation of this intermediate: the spectral signature of partially decomposed LiBH4 closely matched the closo-borane reference compound K2B12H12, confirming [B12H12]2− as a genuine intermediate species rather than earlier-hypothesized substoichiometric “LiBH3/LiBH2/LiBH” phases, while MQMAS NMR further isolated a distinct spectral component matching amorphous elemental boron as the terminal product [9].
The chemical origin of this intermediate was clarified by Friedrichs, Remhof, Hwang and Züttel [7], who demonstrated via in-situ XRD and NMR that Li2B12H12 (together with an amorphous Li2B10H10 phase) forms when LiBH4 reacts with gaseous diborane released during early decomposition at 200˚C—a self-reinforcing mechanism whereby the closo-borate intermediate delays release of free elemental boron until it itself decomposes at higher temperature, a behavior the authors suggest generalizes to other borohydrides [7]. Yan et al. [10] subsequently showed that the balance between the direct decomposition pathway and the indirect Li2B12H12 pathway is jointly controlled by temperature and hydrogen back-pressure: at low pressure (0.1 - 10 bar, 800 - 873 K) both pathways proceed concurrently, whereas elevated H2 pressure (50 bar at 873 K, or 10 bar at 700 K) forces decomposition exclusively through Li2B12H12 [10]. Pitt, Paskevicius, Brown, Sheppard and Buckley [11] mapped this system comprehensively up to 600˚C using DTA/TGA, TPD-MS, XRD, and FTIR: isothermal desorption above the melting point first yields LiH plus a γ-Li2B12H12 polymorph, which converts to a substoichiometric Li2B12H12−x phase persisting largely unchanged up to 600˚C under vacuum with no further LiH formation; only by 650˚C does this residual closo-borate finally decompose, releasing amorphous boron (with some dissolved Li) as the true end product, with prolonged isothermal treatment also generating Li-B alloy side phases. Complementary in-situ evidence comes from Reed and Book [12], whose temperature-resolved Raman spectroscopy (80˚C - 600˚C, flowing argon) directly tracked LiBH4’s phase transition, melting, and decomposition in real time: Li2B12H12 was observed forming at ~350˚C - 356˚C, with amorphous boron appearing shortly after at ~380˚C - 385˚C, and both phases persisted within the melt and remained detectable after cooling to room temperature.
At the atomistic level, Hoang and Van de Walle [13] proposed via first-principles DFT calculations of native point defects in LiBH4 that decomposition is mediated by mass transport of specific defect species: negatively charged hydrogen interstitials ( H i − ) act as the nucleation site for LiH formation and constitute the rate-limiting diffusing species with a self-diffusion activation energy of 1.65 eV, while boron-vacancy defects (V+_BH4, V0_BH3) mediate transport of (BH4)−/BH3 units toward the surface, where elemental boron is ultimately released, and highly mobile lithium vacancies and interstitials (formation energies as low as 0.50 eV, migration barriers ~0.29 - 0.30 eV) maintain local charge neutrality throughout. Beyond mechanistic characterization, recent patent literature by Sanchez et al. [14] has addressed the practical challenge of converting borohydride decomposition into genuinely high-purity, high-surface-area boron: heating an alkali borohydride together with a molten inorganic alkali halide salt (80 - 99.9 wt%) under inert atmosphere at 600˚C - 1000˚C—sufficient both to decompose the borohydride and to keep the salt fully molten so as to disperse the resulting boron particles—yields boron nanoparticle aggregates ≤25 nm (preferably 2 - 10 nm) at >90 - 98 mol% purity and specific surface areas up to ~900 m2·g−1, a substantial improvement over the ~10 - 20 m2·g−1 characteristic of conventional micrometric amorphous boron produced by magnesiothermic reduction of boron oxide.
It should be noted that the thermal decomposition of LiBH4 has been extensively studied in the context of hydrogen storage, where the primary objective is reversible hydrogen release and uptake. In those studies, the Li2B12H12 intermediate is often regarded as an obstacle because it raises the decomposition temperature and limits reversibility [7][10]. In contrast, the present work focuses on the same decomposition process with a different objective: maximizing the yield of ultrapure amorphous elemental boron from LiBH4, which is itself synthesized from LiH and BF3·Et2O. This shift in perspective means that the Li2B12H12 intermediate, rather than being an impediment, is simply a necessary stage on the pathway to elemental boron, and the reaction conditions are optimized accordingly.
2. Experimental2.1. Materials and Reagent Preparation Reagents
All reagents used in this work were of analytical or synthesis grade and were purchased from established international chemical suppliers. The complete list of materials, together with their formulae, purities, suppliers, and catalogue numbers, is given below.
Lithium hydride (LiH) was obtained from Sigma-Aldrich (Merck KGaA, Germany, catalogue no. 201049) as a powder of ≥97% purity with a particle size of ≤0.2 mm. Boron trifluoride (BF3) gas of ≥99.5% purity, supplied anhydrous in a lecture bottle, was also purchased from Sigma-Aldrich (catalogue no. 292532).
The ethereal solvents used as reaction media were all sourced from Sigma-Aldrich (Merck KGaA, Germany): anhydrous diethyl ether ((C2H5)2O) of ≥99.7% purity, containing BHT as stabilizer (catalogue no. 296082); anhydrous tetrahydrofuran (C4H8O) of ≥99.9% purity, inhibitor-free (catalogue no. 401757); and anhydrous 1,2-dimethoxyethane (C4H10O2) of ≥99.5% purity, inhibitor-free (catalogue no. 259527).
Inert gases used throughout the work were high-purity argon (Ar, 99.999% grade 5.0, O2 and H2O < 3 ppm) and high-purity nitrogen (N2, 99.999% grade 5.0), both supplied by Air Liquide (France) and Messer.
Because lithium hydride, boron trifluoride, and lithium borohydride all react rapidly and exothermically with atmospheric moisture and oxygen, all reagents were opened, stored, and handled exclusively inside an argon-filled glove box (MBraun LABstar or equivalent) in which O2 and H2O levels were continuously monitored and maintained below 1 ppm. Solid reagents were transferred to and from the glove box in tightly sealed Schlenk vessels flushed with argon.
Drying and Deoxygenation of Solvents
The ethereal solvents used both as reaction media (diethyl ether, THF, 1,2-dimethoxyethane) and as scrubbing media required rigorous removal of water and dissolved oxygen prior to use. Even trace moisture hydrolyzes BF3, releases HF, deactivates LiH, and destroys any LiBH4 formed; trace oxygen leads to formation of oxidized boron species (BO3, BO4) detectable by 11B NMR. The solvents were therefore dried by a two-stage procedure.
In the first stage, the commercial anhydrous solvent (typically supplied at <50 ppm H2O) was pre-dried by standing over freshly activated 4 Å molecular sieves for a minimum of 48 h. The molecular sieves themselves had been activated by heating at 300˚C under dynamic vacuum (≤10−2 mbar) for 12 h and then cooled and stored under argon before use. This step reduced the residual water content to <10 ppm as verified by Karl Fischer titration (Metrohm 831 KF Coulometer).
In the second stage, the pre-dried solvent was transferred to a solvent still under argon and refluxed over freshly cut sodium metal in the presence of benzophenone as an indicator. The characteristic deep-blue-to-purple color of the sodium benzophenone ketyl radical confirmed the absence of both moisture and oxygen. Solvents were freshly distilled from the still directly into oven-dried Schlenk flasks under argon flow immediately before use; solvent standing in storage for more than 48 h after distillation was re-distilled. For diethyl ether specifically, dissolved peroxides were first removed by passage over an activated basic alumina column prior to sodium/benzophenone distillation, since peroxides in ether pose a serious explosion hazard and interfere with subsequent BF3 complexation.
For 1,2-dimethoxyethane, which is less amenable to sodium/benzophenone drying because of its higher boiling point (85˚C) and coordinating character, the procedure was modified: after pre-drying over 4 Å sieves, the solvent was refluxed for 6 h over CaH2 under argon, then distilled onto fresh molecular sieves and stored under argon in a Schlenk flask sealed with a Teflon stopcock.
Although BF3·Et2O is commercially available, in-house synthesis affords higher purity and allows precise control of the BF3/Et2O molar ratio, which is critical for the subsequent LiBH4 synthesis step. The complex was prepared by direct complexation of gaseous BF3 with rigorously dried diethyl ether according to:
BF3 (g) + Et2O (l) → BF3·Et2O (3)
The reaction was carried out in a jacketed glass reactor (500 mL, PTFE-sealed) fitted with a mechanical stirrer, a gas-inlet dip tube reaching below the liquid surface, a reflux condenser, and an exhaust line vented through a scrubber filled with 5 M aqueous NaOH to trap any unreacted BF3. The entire apparatus was oven-dried at 130˚C for 12 h, assembled hot, and purged with high-purity argon for at least 30 min before use.
Freshly distilled, anhydrous diethyl ether (200 mL, ~1.9 mol) was charged to the reactor under argon flow and cooled to 0˚C - 5˚C in an ice/water bath. Gaseous BF3 was then introduced slowly (~50 mL·min−1, measured by a calibrated mass-flow controller) through the dip tube under vigorous stirring (400 rpm). The reaction is strongly exothermic; the jacket temperature was maintained at 0˚C - 10˚C throughout the addition to prevent evaporation of the ether and to keep the growing BF3·Et2O complex in solution. A slight molar excess of BF3 over Et2O (BF3:Et2O ≈ 1.05:1) was used to ensure complete conversion.
Excess BF3 was then purged from the reactor with dry argon for 30 min. The crude BF3·Et2O was fractionally distilled under argon through a short Vigreux column, collecting the main fraction at 125˚C - 127˚C (lit. b.p. 126˚C) as a colorless-to-pale-yellow liquid. The distilled complex was analyzed by 11B NMR (single sharp signal at +0.5 ppm relative to BF3·Et2O reference) and by density measurement (1.125 g·mL−1 at 25˚C), both consistent with high-purity BF3·Et2O. Because the complex hydrolyzes rapidly in humid air with release of HF, the product was stored in a tightly sealed, PTFE-stoppered Schlenk flask under argon at 4˚C, away from light, and used within four weeks of preparation.
2.3. Synthesis of Lithium Borohydride (LiBH
<sub>4</sub>
)
LiBH4 was prepared by the reaction of finely divided, pre-dried LiH with BF3·Et2O in a higher-boiling ethereal solvent (THF), following the optimized process reported by Hauk and Wietelmann [5]:
4LiH + BF3·Et2O → LiBH4 + 3LiF↓ + Et2O↑
The reaction was carried out in a 500 mL jacketed glass reactor fitted with a mechanical stirrer (300 - 400 rpm), a pressure-equalizing dropping funnel, a reflux condenser, and an argon inlet/outlet system with an exhaust scrubber filled with dry acetone to trap any diborane emissions. All glassware was oven-dried at 130˚C for 12 h and assembled hot under argon flow.
Freshly purified LiH powder (42.4 g, 5.34 mol, LiH:BF3 = 4.73:1) was suspended in 300 g of freshly distilled anhydrous THF inside the glove box and transferred to the reactor under argon. A small seed charge of commercial LiBH4 (1.1 g, ~4 mol% of the BF3 to be added) was added to reliably initiate the exothermic reaction. The BF3·Et2O (158.1 g, 1.13 mol total, ≈48% BF3 w/w) was diluted with anhydrous THF (1:1 v/v) in the dropping funnel and metered dropwise (1 - 2 drops·s−1) into the stirred LiH suspension. The internal temperature was initially raised to 53˚C to trigger the reaction and then maintained at 50˚C - 55˚C by jacket cooling (jacket at 45˚C) throughout the ~90 min addition. Onset of reaction was confirmed by a rapid rise in internal temperature (53˚C → 59˚C within 5 min) and a corresponding increase in solution conductivity (from 0.03 to 0.10 mS·cm−1).
After the BF3·Et2O addition was complete, the mixture was maintained under reflux (66˚C) for a further 90 min to drive the reaction to completion, monitored by the cessation of H2 evolution and stabilization of internal temperature and conductivity. The reaction mixture was then cooled to room temperature and filtered under a slight positive argon pressure through a Schlenk glass frit (porosity 4) to remove the LiF precipitate. The filter cake was washed twice with 100 mL portions of anhydrous THF to recover any entrained LiBH4. The combined THF filtrates yielded 390 g of a clear, colorless LiBH4 solution (5.9 wt% LiBH4, corresponding to 22.9 g isolated, 93% yield based on BF3).
For isolation of solid LiBH4, the filtered solution was concentrated on a rotary evaporator (30˚C - 40˚C, 10 - 20 mbar) under an argon blanket, and the resulting white crystalline solid was dried under high vacuum (10−2 mbar) at 60˚C - 80˚C for 4 - 6 h to remove coordinated THF. The product was analyzed by 11B NMR, showing a single dominant signal at −41.3 ppm (>99% of total signal, characteristic of [BH4]−) and a very minor signal at −30 ppm (~0.5%, trace secondary product), with fluoride content below detection limit (<0.005 mmol·g−1) and no boron detected in the acetone scrubber (<0.005 mmol·g−1, confirming negligible diborane emission). The isolated LiBH4 was stored under argon in a sealed Schlenk flask at 4˚C in the dark.
2.3.1. Effect of LiH:BF3 Molar Ratio on LiBH4 Yield
To establish the optimum stoichiometry for LiBH4 formation, a series of eight parallel reactions was carried out at fixed reaction conditions (THF as solvent, addition temperature 50˚C - 55˚C, reflux time 90 min, BF3·Et2O metered into LiH suspension over 90 min, 4 mol% LiBH4 seed charge), varying only the LiH:BF3 molar ratio from a substoichiometric 3.5:1 up to a large excess of 5.0:1. In each experiment, the amount of BF3·Et2O was held constant at 1.13 mol, and the mass of LiH was adjusted accordingly. LiBH4 yield was determined gravimetrically after filtration, solvent removal, and vacuum drying, and cross-checked by quantitative 11B NMR against a calibrated internal standard. The results are summarized in Table 1.
The yield increased markedly between 4.1:1 (46% ± 2%) and 4.3:1 (71% ± 2%), consistent with the >4.1:1 threshold specified in the Hauk & Wietelmann patent [5]. At sub stoichiometric and near-stoichiometric ratios (Entries 1 - 4, 3.5:1 to 4.1:1), LiBH4 yields remained low (32% - 46%), and the 11B NMR spectra of the crude product were heavily contaminated with LiBF4 (−1 ppm), LiB2H7 (−25 ppm), and BH3 residues (−15 to +5 ppm). This is consistent with Fedneva’s original observation that, at ratios of 4:1 or below, the good solubility of LiF in the reaction medium favors formation of LiBF4 as a competing thermodynamic sink, effectively consuming boron before it can be reduced to [BH4]−. At 4.7:1 the yield reached a plateau (93% ± 2%); a further increase to 5.0:1 gave no significant improvement (94% ± 2%, p > 0.05). Product purity, as measured by 11B NMR, followed the same trend, reaching 99% for entries 7 and 8, with essentially no detectable LiBF4 or diborane-related species. A ratio of 4.7:1 was therefore selected as the lowest ratio providing the maximum yield and purity.
Table 1. Influence of LiH:BF3 molar ratio on LiBH4 yield and product purity. Conditions: 1.13 mol BF3·Et2O in THF, addition at 50˚C - 55˚C, 90 min addition + 90 min reflux, 4 mol% LiBH4 seed.
Entry
LiH:BF
3
(mol/mol)
LiH mass (g)
LiBH
4
yield (%)*
Purity by
11
B NMR (%)
Main byproducts detected
1
3.5:1
31.4
32
61
LiBF
4
, LiB
2
H
7
, BH
3
2
3.8:1
34.1
41
68
LiBF
4
, LiB
2
H
7
, BH
3
3
4.0:1
35.9
44
72
LiB
2
H
7
, BH
3
, traces LiBF
4
4
4.1:1
36.8
46
74
LiB
2
H
7
, BH
3
5
4.3:1
38.6
71
88
LiB
2
H
7
(traces)
6
4.5:1
40.4
87
94
Traces LiB
2
H
7
7
4.7:1
42.2
93
99
None detected
8
5.0:1
44.9
94
99
None detected
*Yield based on BF3 charged. All experiments were performed in triplicate (n = 3); values are given as mean ± SD, with SD ≤ 2%.
2.3.2. Effect of Reaction Temperature on LiBH4 Yield
Having fixed the LiH:BF3 ratio at the optimum value of 4.7:1, a second series of five reactions was performed to establish the influence of reaction temperature on LiBH4 yield and purity. All other conditions were held constant (THF as solvent, LiH:BF3 = 4.7:1, 90 min BF3·Et2O addition, 90 min post-addition reflux, 4 mol% LiBH4 seed). The internal reaction temperature during the BF3·Et2O addition and subsequent hold was varied from 25˚C (well below the THF reflux temperature) up to 66˚C (reflux of THF). Results are summarized in Table 2.
The temperature series shows a clear dependence of yield on reaction temperature. At 25˚C (Entry 9), the reaction between LiH and BF3·Et2O proceeded far too slowly: internal temperature and conductivity measurements indicated that only ~50% of the BF3 had reacted by the end of the 90 min hold, and analysis of the isolated product showed substantial residual LiBF4, reflecting incomplete reduction. Raising the temperature to 35˚C (Entry 10) accelerated conversion significantly but not sufficiently for complete reaction within the given hold time, and traces of LiB2H7 remained detectable. At 45˚C (Entry 11) the reaction proceeded smoothly to 85% ± 2% yield with 94% purity. At 55˚C (Entry 12) the yield increased significantly to 93% ± 2% (p < 0.01), with 99% 11B NMR purity and no detectable fluoride or diborane species—consistent with the reaction being fast enough to consume all BF3 within the addition period, while remaining well below the reflux temperature of THF, so that diborane (which is only sparingly soluble in THF) remains dissolved and reacts further with excess LiH rather than escaping.
Table 2. Influence of reaction temperature on LiBH4 yield and product purity. Conditions: LiH:BF3 = 4.7:1, THF solvent, 90 min addition + 90 min hold, 4 mol% LiBH4 seed.
Entry
Reaction temperature (˚C)
LiBH
4
yield (%)*
Purity by
11
B NMR (%)
Notes/observations
9
25
42
74
Very slow reaction; incomplete BF
3
consumption after hold; residual LiBF
4
detected
10
35
68
87
Reaction proceeds smoothly but slowly; some LiB
2
H
7
still present
11
45
85
94
Good conversion; only trace LiB
2
H
7
12
55
93
99
Optimum: full conversion, clean
11
B NMR, no fluoride
13
66 (reflux)
89
96
Slight loss to volatile diborane observed in scrubber (0.04 mmol B·g
−1
)
*Yield based on BF3 charged. All experiments were performed in triplicate (n = 3); values are given as mean ± SD, with SD ≤ 2%.
At the highest tested temperature of 66˚C (Entry 13, reflux), the yield was lower (89% ± 2%). Although this difference is close to the experimental uncertainty (p ≈ 0.07), a small but detectable amount of boron (0.04 mmol B·g−1 solution equivalent) was captured in the acetone scrubber only at this temperature, indicating partial loss of the boron budget as gaseous diborane escaping the reaction mixture before it could react with the surrounding LiH. Additionally, 11B NMR of the crude product at this temperature showed a small peak at −25 ppm attributable to trace LiB2H7, consistent with incomplete second-step reduction. A reaction temperature of 50˚C - 55˚C was therefore selected for all subsequent syntheses: it is sufficiently high to ensure full kinetic conversion of BF3·Et2O within the addition/hold window, but low enough to retain any transiently formed diborane in solution long enough for its complete conversion to LiBH4.
2.4. Thermal Decomposition of LiBH
<sub>4</sub>
to Amorphous Elemental Boron Apparatus
Thermal decomposition of the isolated LiBH4 was carried out in a horizontal high-temperature tube furnace (Carbolite Gero model STF 16/450, Carbolite Gero Ltd., Hope Valley, UK) equipped with a maximum operating temperature of 1600˚C and a heated zone of 450 mm length. The furnace was fitted with a dense recrystallized alumina (Al2O3, ≥99.7% purity) work tube (outer diameter 60 mm, inner diameter 50 mm, total length 900 mm), sealed at both ends with water-cooled stainless-steel vacuum flanges equipped with Viton O-rings and Swagelok gas-inlet and gas-outlet ports.
Temperature control and programming were provided by a Eurotherm 3216P1 PID controller integrated with the furnace, capable of programmable multi-segment heating and cooling ramps with a control accuracy of ±1˚C. The sample temperature was independently monitored by a Type-S (Pt/Pt-10%Rh) thermocouple positioned in direct contact with the sample crucible via a closed-end alumina thermowell inserted through the outlet flange, with temperature data logged at 5-second intervals to a PC.
The gas-handling system consisted of: (i) an inlet line supplying high-purity argon (99.999%, Messer, Georgia) passed sequentially through an oxygen trap (Agilent OT3-4, <15 ppb O2 output) and a moisture trap (Agilent MT200-4, <15 ppb H2O output) to remove residual contaminants; (ii) a calibrated mass-flow controller regulating argon flow through the tube; and (iii) an outlet line vented through a bubbler filled with silicone oil (to prevent back-diffusion of air) followed by a scrubber containing 5 M aqueous NaOH to neutralize any escaping acidic gases. A turbomolecular pumping station (base pressure ~10−5 mbar) was connected to the outlet flange through a UHV-compatible bellows valve, allowing evacuation of the tube prior to argon backfilling. Pressure was monitored by a combined gauge mounted on the outlet flange.
The sample crucibles were of stainless steel (grade 316L, wall thickness 1 mm, internal dimensions 60 × 20 × 15 mm).
2.4.1. Sample Preparation and Loading
All sample handling was carried out inside the argon-filled glove box. A weighed portion of freshly prepared LiBH4 (typically 5.0 g, corresponding to a theoretical elemental boron yield of 2.48 g assuming complete decomposition according to Equation (3) below) was transferred as a fine powder into a pre-cleaned, oven-dried (130˚C, 12 h) 316L stainless-steel crucible. The crucible was loaded no more than half-full to allow for the volume expansion accompanying melting (~15%) and to provide sufficient headspace for the escape of evolved H2 gas.
The loaded crucible was placed inside a purpose-built stainless-steel transfer container fitted with a gas-tight O-ring seal and sealed under argon. The transfer container was removed from the glove box and connected directly to a modified port on the furnace outlet flange, and the crucible was slid into the center of the alumina work tube (i.e., the isothermal zone) via a stainless-steel push rod under a continuous positive argon flow (100 sccm) that prevented ingress of air. Once the crucible was correctly positioned within ±5 mm of the thermocouple tip, the transfer port was closed and the furnace tube sealed.
2.4.2. Decomposition Procedure
The overall decomposition reaction targeted is:
2LiBH4 → 2LiH + 2B + 3H2↑ (4)
but proceeds in practice through the intermediate Li2B12H12 borate phase:
12LiBH4 → Li2B12H12 + 10LiH + 13H2↑ (5)
Li2B12H12 → 2LiH + 12B + 5H2↑ (6)
so that reaching 600˚C - 650˚C is essential to fully decompose the persistent Li2B12H12 intermediate and release amorphous elemental boron.
The furnace tube was first evacuated to a base pressure of ~10−4 mbar using the turbomolecular pump, held for 30 min to remove any residual atmosphere, and then backfilled with high-purity argon to slightly above atmospheric pressure (1.05 bar). This evacuation-backfill cycle was repeated three times to ensure a rigorously oxygen- and moisture-free atmosphere in the work tube. During the entire decomposition run, a continuous slow argon flow (100 sccm) was maintained through the tube to sweep out evolved H2 and any volatile byproducts (traces of diborane), keeping the interior atmosphere inert and preventing back-diffusion of air through the outlet.
The decomposition thermal program consisted of five segments:
Segment 1—Warm-up (25˚C→ 200˚C, 5˚C·min−1, 35 min). The furnace was heated at a moderate ramp rate to dry the sample and desorb any weakly adsorbed species without inducing premature decomposition.Segment 2—Pre-melt hold (200˚C, 30 min). An isothermal hold was applied to allow any final moisture desorption and to equilibrate the temperature across the sample.Segment 3—Melting and initial decomposition (200˚C→ 350˚C, 3˚C·min−1, 50 min). The slow ramp rate through the melting region (LiBH4 m.p. ≈ 280˚C) allowed controlled melting without excessive foaming from the accompanying H2 evolution. Onset of hydrogen release was indicated qualitatively by a rise in outlet pressure (~1.1 - 1.3 bar) despite the continuous argon purge, consistent with formation of the Li2B12H12 intermediate and initial amorphous boron nuclei according to Equation (5).Segment 4—High-temperature decomposition (350˚C→ 650˚C, 5˚C·min−1, 60 min). The furnace was ramped to the target dwell temperature of 650˚C, chosen to lie above the reported ~600˚C stability limit of Li2B12H12−x, ensuring full decomposition of the closo-borate intermediate to elemental boron and LiH per Equation (5).Segment 5—Isothermal dwell (650˚C, 4 h). The sample was held isothermally at 650˚C for 4 h to drive the reaction to completion. The outlet pressure was used only for qualitative monitoring of hydrogen evolution; no calibration to H2 flow was performed. The fixed 4 h dwell was chosen to exceed the point at which no further pressure change was observed over a 30 min window.
After the dwell, the furnace was cooled slowly to room temperature at 5˚C·min−1 under continuous argon flow to prevent thermal shock to the alumina tube and to avoid re-oxidation of any exposed boron surfaces. Once the internal temperature had dropped below 40˚C, the outlet flange was opened under argon flow, the sample crucible was transferred back to the sealed transfer container while still under argon protection, and the container was returned to the glove box for product recovery.
2.5. Purification of Amorphous Elemental Boron Rationale
The dark grey/black solid recovered from the tube furnace after thermal decomposition of LiBH4 is not pure boron. It consists of amorphous elemental boron intimately mixed with several coproducts and residual intermediates that must be systematically removed to reach the target purity of >98.5%. The principal contaminants are:
Lithium hydride (LiH)—the stoichiometric coproduct of LiBH4 decomposition (Equation (3)), present in molar equivalence to boron;Residual Li2B12H12−x—small amounts of the persistent substoichiometric closo-borate intermediate that may survive if any portion of the sample experienced local under-heating;Trace lithium metal—formed by partial thermal decomposition of LiH at the highest dwell temperatures;Surface oxide/hydroxide layer (B2O3, B(OH)3)—formed rapidly on freshly exposed boron surfaces by reaction with atmospheric moisture and oxygen during any brief handling excursion outside the glove box;Metallic contaminants from the crucible—traces of Fe, Cr, and Ni that may have leached from the 316L stainless-steel crucible during the isothermal dwell at 650˚C.
Each contaminant class requires a specific purification strategy exploiting differences in solubility, reactivity, and chemical stability between it and elemental boron. The purification protocol described below proceeds in five sequential stages, each targeted at removing a specific contaminant class while preserving the amorphous boron matrix intact.
Stage 1—Controlled Ethanolysis of Lithium Hydride
The most abundant contaminant, LiH, is highly reactive and reacts violently with water according to:
LiH + H2O → LiOH + H2↑(7)
Direct addition of water to the crude powder would therefore cause vigorous, potentially incendiary H2 evolution, thermal shock, and mechanical disruption of the amorphous boron particles. To avoid this, LiH is first hydrolyzed in a controlled manner using absolute ethanol as a milder proton source:
LiH + C2H5OH → LiOC2H5 + H2↑ (8)
The reaction of LiH with ethanol proceeds smoothly at 0˚C, with H2 evolution rate low enough to be safely managed by an argon purge and a mineral oil bubbler.
Stage 2—Aqueous Quench and Acid Wash
Once the LiH was fully converted to soluble lithium ethoxide, distilled water (100 mL) was added dropwise to the suspension over 30 min, again with cooling in an ice/water bath. This hydrolyzed the ethoxide to lithium hydroxide and simultaneously dissolved any residual water-soluble Li-containing species. Critically, Li2B12H12 and its substoichiometric derivative Li2B12H12−x are highly water-soluble (>50 g/100 mL H2O at 25˚C) and pass into the aqueous phase at this stage, while amorphous elemental boron is completely insoluble in water and remains suspended as a dark solid.
Following aqueous quenching, dilute hydrochloric acid (2 M HCl, 100 mL) was added slowly to the suspension with continued stirring, bringing the medium to pH ≈ 1. The acidic conditions dissolve:
residual LiOH, forming soluble LiCl;any surface B2O3/B(OH)3 oxide layer, converting it to soluble boric acid (H3BO3);traces of Fe, Cr, and Ni leached from the crucible, converting them to soluble metal chlorides;any remaining Li2B12H12−x, which decomposes acidolytically.
Amorphous boron itself does not react with dilute HCl at room temperature and survives this treatment unchanged. The acid-treated suspension was stirred at room temperature for a further 3 h to allow complete dissolution of the contaminants.
Stage 3—Filtration and Neutral Washing
The acid-treated suspension was filtered under a slight positive argon pressure through a Büchner funnel fitted with a POR-4 sintered glass frit (porosity 10 - 16 μm), retaining the fine amorphous boron powder as a dark cake while allowing all dissolved contaminants to pass through with the filtrate. The filter cake was washed sequentially: With distilled water (5 × 50 mL portions), until the filtrate pH reached neutral (pH 6.5 - 7.0), with absolute ethanol (3 × 30 mL portions), to displace residual water from the porous boron structure and to accelerate subsequent drying, with acetone (3 × 30 mL portions), to further reduce residual solvent content and prevent aggregation of the fine particles during drying.
Stage 4—Vacuum Drying
The washed cake was transferred, still on the sintered glass filter, into a vacuum oven and dried at 80˚C - 100˚C under reduced pressure (<10 mbar) for 6 h.
3. Result and Discussion
3.1. Thermal Decomposition of LiBH
<sub>4</sub>
—Derivatographic Analysis
Simultaneous TG/DTA/DSC (Netzsch STA 449 F5 Jupiter, alumina crucible, argon flow 60 mL·min−1, heating rate 5˚C·min−1, from 30˚C to 800˚C) of a 25 mg sample of the isolated LiBH4 produced a derivatogram (Figure 1) showing four distinct thermal events:
1) ~116˚C—a small, sharp endothermic peak with negligible mass loss (<0.1 wt%), corresponding to the orthorhombic → hexagonal polymorphic transition of solid LiBH4.
2) ~278˚C—a sharp endothermic peak with a small associated mass loss of ~0.4 wt%, corresponding to the melting of LiBH4 accompanied by a minor initial hydrogen desorption from the liquid.
3) 320˚C - 450˚C (peak at ~385˚C)—a broad endothermic feature with major mass loss of ~7.8 wt%, corresponding to the first significant decomposition stage in which LiBH4 decomposes to Li2B12H12 and LiH (Equation (4)) with release of H2. The mass loss agrees well with the theoretical value of 8.0 wt% expected for stoichiometric conversion of 12LiBH4 to Li2B12H12 + 10LiH + 13H2.
4) 500˚C - 680˚C (peak at ~600˚C)—a second broad endothermic feature with a further mass loss of ~4.6 wt%, corresponding to decomposition of the Li2B12H12−x intermediate to LiH + amorphous B + H2 (Equation (5)). Beyond 680˚C, no further mass loss was detected up to 800˚C, confirming that the LiH coproduct is stable in this temperature range, consistent with its literature decomposition temperature of ~900˚C.
The total measured mass loss up to 680˚C was 12.8 wt%, in reasonable agreement with the theoretical value of 13.8 wt% for full conversion of LiBH4 to LiH + B + (3/2) H2 (Equation (3)), with the ~1 wt% shortfall attributable to a small loss of boron budget as trace diborane during the intermediate stage, consistent with the observations of Friedrichs et al. [7] and of Reed & Book [12]. The clear separation between the two hydrogen-release events in the derivatogram (peaks at ~385˚C and ~600˚C) is consistent with the two-step decomposition pathway via a persistent Li2B12H12 intermediate, as proposed by Orimo et al. [8] and mapped in detail by Pitt et al. [11].
Figure 1. Derivatogram (TG/DTA/DSC) of the thermal decomposition of LiBH4.
3.2. Confirmation of Amorphous Boron Formation
The purified end-product of the tube-furnace decomposition was extensively characterized to confirm both its amorphous nature and its elemental identity.
XRD of the purified powder (Bruker D8 Advance, Cu Kα, 2θ = 10˚ - 80˚, air-tight sample holder) showed no crystalline reflections attributable to LiH, LiF, Li2B12H12, LiBH4, or any crystalline polymorph of boron. Instead, only two very broad diffuse halos were observed (Figure 2), centered at 2θ ≈ 20˚ and ~35˚ (FWHM ~10˚ and ~15˚, respectively), characteristic of the short-range order of amorphous elemental boron. This pattern matched closely with the reference XRD trace of commercial amorphous boron (ABCR, 95% - 97% purity) and with the pattern reported by Sanchez et al. [14] for their nanostructured amorphous boron.
Figure 2. XRD pattern of the purified amorphous boron powder, showing broad diffuse halos characteristic of amorphous elemental boron.
Figure 3. SEM image of the purified amorphous boron powder: primary particles of 40 - 100 nm aggregating into 200 - 500 nm clusters.
SEM-EDXrevealed a nanoparticulate morphology (Figure 3) consisting of aggregates of primary particles in the 40 - 100 nm range, with the larger aggregates reaching 200 - 500 nm. EDX elemental mapping showed a uniform boron distribution with only trace oxygen signal (~1.5 at% O), and no detectable Li, F, or transition metal signals above the noise floor (~0.1 at%). This indicates both the compositional purity of the product and the complete removal of the LiF/LiCl residuals from the acid-washing stages.
ICP-OES elemental assay after microwave-assisted acid digestion (concentrated HNO3/HF, CEM MARS 6) gave an elemental boron content of 98.7% ± 0.2%, with Li = 32 ppm, F below detection limit (<10 ppm), Fe = 45 ppm, Cr = 15 ppm, Ni = 8 ppm, and Al below detection limit (<10 ppm). The total metallic contamination remained below 100 ppm, meeting the specification for high-purity amorphous boron. The boron content (98.7 wt%, determined by [method]) reflects the bulk composition. The oxygen detected by EDX (~1.5 at%) arises from a thin B2O3 surface layer formed on air exposure of the high-surface-area boron during sample handling, which EDX, as a surface-sensitive method, overrepresents. The remaining ~1.3 wt% is attributed mainly to oxygen, hydrogen, and carbon.
FTIR (PerkinElmer Frontier, ATR mode) of the purified boron powder (Figure 4). showed complete absence of the B-H stretching bands at 2400 - 2600 cm−1 that would signal residual borohydride or closo-borate species, confirming the completeness of the thermal decomposition. Only weak broad absorption near 1350 - 1450 cm−1 was observed, attributable to the trace surface BO3 groups already identified by 11B NMR.
Figure 4. FTIR spectrum of the purified boron powder: no B-H stretch (2400 - 2600 cm−1), weak trace BO3 band (1350 - 1450 cm−1).
4. Conclusions
This work has demonstrated that ultrapure amorphous nanostructured boron can be reliably prepared through a two-stage laboratory-scale route based on well-established but often overlooked chemistry: synthesis of lithium borohydride from lithium hydride and boron trifluoride diethyl etherate in tetrahydrofuran, followed by controlled thermal decomposition of the isolated LiBH4 in a horizontal tube furnace under argon at 650˚C, and a staged aqueous purification (ethanolysis of LiH followed by dilute HCl washing) to remove all residual lithium- and halide-containing species.
Systematic variation of the LiH:BF3 molar ratio and reaction temperature identified an optimum at LiH:BF3 = 4.7:1 and 50˚C - 55˚C, delivering LiBH4 in 93% isolated yield and >99% 11B NMR purity, in close quantitative agreement with the industrial process patented by Hauk & Wietelmann. Thermal decomposition proceeded in two steps, with sequential mass losses at ~385˚C and ~600˚C matching the theoretical values for Equations (4) and (5), consistent with the pathway via a Li2B12H12 intermediate reported in the literature [8][11]. The final purified product was characterized as amorphous elemental boron of 98.7% purity, with primary particle size ~60 nm, specific surface area 45 m2·g−1.
The two-stage LiH/BF3-based route therefore represents a practical, scalable, and reproducible pathway to nanostructured amorphous boron from readily available industrial reagents—well suited to producing precursor material for metal borides, high-energy-density additives, and other advanced boron-based applications where reactivity and purity outweigh the need for the ultrafine morphologies achievable only by molten-salt-mediated methods.
Author Contributions
Paata Nikoleishvili: Conceptualization, Methodology, Writing - original draft, Writing - review & editing—formulation of the main idea of the work and the research strategy; preparation of the manuscript.
Gigla Tsurtsumia: Supervision, Project administration, Investigation—supervision and execution of the project.
Tina Lezhava: Investigation—study of the lithium borohydride synthesis proceses.
Nana Koiava: Investigation—study of the lithium borohydride decomposition processes.
All authors have read and agreed to the published version of the manuscript.
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