Metal boranes: Progress and applications

Metal boranes: Progress and applications

G Model ARTICLE IN PRESS CCR-112165; No. of Pages 11 Coordination Chemistry Reviews xxx (2015) xxx–xxx Contents lists available at ScienceDirect ...

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G Model

ARTICLE IN PRESS

CCR-112165; No. of Pages 11

Coordination Chemistry Reviews xxx (2015) xxx–xxx

Contents lists available at ScienceDirect

Coordination Chemistry Reviews journal homepage: www.elsevier.com/locate/ccr

Review

Metal boranes: Progress and applications Bjarne R. S. Hansen a , Mark Paskevicius a , Hai-Wen Li b,c , Etsuo Akiba b,c,d , Torben R. Jensen a,∗ a Center for Materials Crystallography, Interdisciplinary Nanoscience Center (iNANO), and Department of Chemistry, Aarhus University, Langelandsgade 140, DK-8000 Aarhus C, Denmark b International Research Center for Hydrogen Energy, Kyushu University, 744 Motooka, Nishi-ku, Fukuoka 819-0395, Japan c WPI International Institute for Carbon-Neutral Energy Research (WPI-I2CNER), Kyushu University, 744 Motooka, Nishi-ku, Fukuoka 819-0395, Japan d Department of Mechanical Engineering, Faculty of Engineering, Kyushu University, Fukuoka 819-0395, Japan

Contents 1. 2.

3. 4.

5.

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 Synthesis of metal boranes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 2.1. Route I: Diborane addition . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 2.2. Route II: Iodine reduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 2.3. Route III: Decaborane based . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 2.4. Route IV: Ion exchange . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 2.5. Route V: Carborane synthesis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 Route VI: Functionalization . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 2.6. Structure and coordination in solid metal boranes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 Properties and applications of metal boranes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 4.1. Ion conductivity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 4.2. Boranes in medicine . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 4.3. Carborane and borane-based polymers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 4.4. Higher boranes in hydrogen storage . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 Acknowledgments . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 00

a r t i c l e

i n f o

Article history: Received 14 October 2015 Received in revised form 2 December 2015 Accepted 3 December 2015 Available online xxx

a b s t r a c t Boron and hydrogen have a rich chemistry that has attracted a significant, but mainly academic, interest during the past century. However, research over the past decades has revealed new applications for metal boranes including their implementation as ‘energy materials’ and as neutron capture agents in cancer treatment. The energy applications involve the use of boron–hydrogen compounds as ion-conductors for batteries, as hydrogen storage materials, or even rocket fuels. The intensive research focus on metal borohydrides in the early 21st century has recently broadened to encompass higher metal boranes such as metal–B10 H10 ’s and B12 H12 ’s. This review summarizes the recent breakthroughs in the area of higher metal boranes in these last few years, in addition to highlighting core research from the mid-20th century. © 2015 Elsevier B.V. All rights reserved.

1. Introduction Boron hydrides, or boranes, have been intensively investigated during the past century due to their great variety of structures,

∗ Corresponding author. Tel.: +45 22721486. E-mail address: [email protected] (T.R. Jensen).

compositions and unusual chemical bonding schemes [1]. Boranes are the fourth most extensive group of hydrides after carbon, phosphorus and silicon hydrides. The study of boranes was initiated by Alfred Stock and later Herman I. Schlesinger and Herbert C. Brown, who pioneered the research field [2–4]. The most simple boron–hydrogen compound, BH3 , is clearly electron deficient and will readily dimerise to diborane, B2 H6 , in an exothermic reaction. To describe diborane, W. N. Lipscomb established a “two

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Fig. 1. Structures of selected arachno-, nido- and closo-borane anions and the year they were first synthesized. Increasing number of boron from left to right. The longest distance between terminal H-atoms is listed above the structure [4,15–20].

electron-three center” model for the boron–hydrogen bond [5,6]. In the “two electron-three center” bond three atomic orbitals may combine to form a bonding and two antibonding orbitals or a bonding, nonbonding and antibonding set, i.e., a bridging hydrogen [7]. Formation of “two electron-three center” bonds completes the electron shell for boron to resemble neon. This type of bond is typical for elements with relatively few electrons, e.g., alane, AlH3 , where aluminum coordinates six hydrogen atoms in the solid state, forming corner sharing octahedra, similar to structures of AlF3 . However, boron–hydrogen bonding is clearly more covalent and can be bridging: B H B, or terminal: B H or B H2 . The molecular boranes, Bm Hn , resemble carbohydrides by having different types of B B bonds, in contrast to alane that is only built from bridging Al H Al. The boranes are usually optically transparent, diamagnetic, molecular compounds, which can be subdivided in different classes with the most well-known being closo-, nido-, and arachno-boranes. The most stable class of anions is the closo-boranes, Bn Hn 2− , n = 6–12, which have closed polyhedral clusters of n boron atoms and only terminal hydrogen atoms (Fig. 1). Thus, closo-borane structures can be considered polyhedral with all corners occupied by boron, whereas nido-, (Latin ‘nidus’, English ‘nest’) and arachnoboranes (Latin ‘arachne’, English ‘spider’) are polyhedral with one or two missing boron atoms in the corners of the polyhedra, respectively [8]. The most well-known nido-boranes, Bn Hn+4 , are B2 H6 and B10 H14 . Nido-boranes also form a series of anions with the general formula, Bn H− , starting with the borohydride anion, BH− , n+3 4 which may also be considered as the Lewis acid–base pair BH3 and H− . The series of arachno-boranes, Bn Hn+6 , can be exemplified − − by B4 H10 and the anions, Bn H− n+5 , i.e., B2 H7 , and B3 H8 . Bridging hydrogen atoms, B H␮ B, tend to be more acidic than terminal atoms and are lost first during the formation of anions. The reactivity tends to increase and the stability decreases in the series, closo- < nido- < arachno-boranes. The chemistry of boranes began close to 110 years ago, when the stoichiometry of the boranes: B2 H6 , B4 H10 , B5 H9 , B5 H11 , B6 H10 and B10 H14 , was discovered between 1904 and 1937 [2,3]. From the late 1940s until roughly 1959, the U.S. and Soviet Union secretly invested in the development of borane-based fuels, mainly for military use. This research was directed towards molecular boranes such as diborane, B2 H6 , pentaborane, B5 H9 , decaborane, B10 H14 and ethyl decaborane “HEF-3”. Boron fuels were believed to be powerful propellants, superior to the available hydrocarbon fuels, with 40–50% more energy released from hydroboron than hydrocarbon [9]. While this development was abandoned by the U.S. military in 1959, it did lead to a significant breakthrough in borane production on a sizeable scale (∼5 large scale plants were built in the U.S.) and accelerated boron chemistry with tremendous speed. During this time (1940–1967), a range of nido-(BH− ), arachno-(B3 H− ) and 4 8 2− 2− closo-boranes (B6 H6 − B12 H12 ) were all synthesized/reported for

the first time (Fig. 1). The size of certain modified anions (para˚ is comparable to C2 B10 (CH3 )12 , 9.9 A˚ and para-C2 B8 H8 Cl2 , 14.4 A) ˚ [10]. Interestingly, the C60 fullerenes show the size of C60 (10.7 A) reorientational motion above −13 ◦ C, similar to closo-boranes in their high temperature polymorphs [11]. Perhaps these two compounds, with comparable size and symmetry (point group Ih ) share other properties. For instance, alkali- and alkaline earth metal C60 complexes display rapid ion conductivity, similar to alkali metal B12 H12 , in part owing to three-dimensional ionic diffusion [12–14]. 2. Synthesis of metal boranes A large variety of synthesis methods have been developed over the last 50–60 years to form metal boranes. The most commonly synthesized metal boranes are the rather stable closo-boranes, Bn Hn2− (n = 10, 12), which are focused on here. Many boranes are highly toxic when inhaled or absorbed through the skin, which has resulted in some cases of poisoning and even death [21]. Certain boranes may also react aggressively in contact with air or moisture. In general, the lower boranes are less stable and more reactive, thus more dangerous to handle. Metal–B12 H12 compounds are considered to be reaction intermediates in the decomposition of many metal borohydrides, e.g., Li2 B12 H12 from LiBH4 [22]. There is a close relationship between metal borohydrides and higher metal boranes, not only in their decomposition but also their synthesis. Hawthorne first syn2− thesized the B12 H12 anion in very low yields in 1960 from a 2-iododecaborane precursor [18]. The synthesis of a metal–B12 H12 compound can be achieved through a number of routes from different boron-rich precursors (see Fig. 2). Generally, high temperatures ∼∼180 ◦ C are required to initiate boron–boron bond building (B H 2− condensation), forming closed B12 H12 icosahedra. Below these temperatures lower metal boranes are formed instead, such as MB11 H14 . 2.1. Route I: Diborane addition The reaction of diborane gas with a metal borohydride results in a series of boron-rich anions, dependent on reaction temperature, which tend to be larger and more stable with increasing reaction temperature. For instance, in 1963, the reaction between B2 H6 and NaBH4 was shown to proceed to Na2 B12 H12 when reacted in triethylamine at 100–180 ◦ C [23]. However, this reaction could also be performed at 25 ◦ C in dimethoxyethane to produce NaB3 H8 , and when the reaction temperature was raised above 50 ◦ C a mixture of NaB11 H14 and Na2 B12 H12 was formed. Recently, solventfree methods have been employed by direct solid–gas reactions between borohydrides and diborane [24,25]. However, high yields of MB12 H12 are only obtained when reactions between MBH4 and

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Fig. 2. Illustration of different metal Bn Hn 2– (n = 6–12) synthesis pathways, mainly B10 H10 2− and B12 H12 2– , as described in the text.

B2 H6 are undertaken within a mechanical mill to improve diborane permeability through the as-formed MB12 H12 layer. Milling under elevated temperatures can also aid in the formation of MB12 H12 .

compound, as illustrated by Reaction (3) and (4) (See Route III.b in Fig. 2): [29] B10 H14 (s) + 2(CH3 )2 S(l) → B10 H12 ·2S(CH3 )2 (s) + H2 (g)

(3)

2.2. Route II: Iodine reduction 2− Reasonable yields of a B12 H12 salt were reported in 1963 through the thermal treatment of NaB3 H8 in diglyme (diethylene glycol dimethyl ether) under reflux [26]. The synthesis can be made cheap and straightforward by starting with NaBH4 and I2 in diglyme as follows: [27,28]

3NaBH4 + I2 → NaB3 H8 + 2NaI + 2H2 (g)

(1)

5NaB3 H8 → Na2 B12 H12 + 3NaBH4 + 8H2 (g)

(2)

The reaction is completed by thermal treatment under reflux. Acid treatment and the addition of triethylamine affords (Et3 NH)2 B12 H12 . It has been highlighted that this route avoids toxic or expensive precursors (such as B2 H6 or B10 H14 ) [27]. 2.3. Route III: Decaborane based Decaborane (B10 H14 ) is an open nido-borane, but is a toxic solid with a low melting point (m.p. = 98.8 ◦ C), causing it to be reason2− ably volatile. The reaction pathway required to synthesize B10 H10 anions from B10 H14 was first reported by Hawthorne in 1959 [19]. The reaction proceeds through two steps, where first a ligand is bound to the open B10 H14 cage, forming B10 H12 (ligand)2 , which 2− ionic can further react with a base to form an equivalent B10 H10

B10 H12 ·2S(CH3 )2 (s) + 2NH3 (g) → (NH4 )2 B10 H10 (s) + 2(CH 3 )2 S(l)

(4)

2− The formation of B12 H12 from a decaborane precursor is also possible. The reaction between B10 H14 and BH4− ions in diglyme was described in 1964 [30]. The diglyme is proposed to take part in the − ion is reaction, which is temperature dependent, whereby a B10 H13

− 2− formed at room temperature, a B11 H14 ion at 90 ◦ C, and a B12 H12 ion at higher temperature. For example, the reaction between NaBH4 and B10 H14 in refluxing diglyme yields the Na2 B12 H12 ·diglyme salt, which can be dissolved in water and dried to yield Na2 B12 H12 . Only recently was this process conducted without the presence of a solvent to yield alkali (Li, Na and K) and alkaline earth (Mg and Ca) M2/n B12 H12 , where n is the charge of the metal [31,32]. Sintering of the borohydride with B10 H14 in a sealed vessel from 200–450 ◦ C yields the respective alkali and alkali earth metal–B12 H12 compound (See Route III.a in Fig. 2). The reaction between decaborane and borane precursors can 2− also result in the formation of the B12 H12 anion. In 1963 Greenwood undertook: [33]

B10 H14 + 2Et3 NH·BH3 → (Et3 NH)2 B12 H12 + 3H2

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The mechanism of formation is suggested to proceed by dissociation of triethylamine-borane, followed by adducting to decaborane and then rapid electrophilic attack of the complex by the BH3 groups (See Route III.c in Fig. 2). 2.4. Route IV: Ion exchange In 1963, Muetterties used a strongly acidic ion exchange column to obtain acidic (H3 O)2 B12 H12 from a Na2 B12 H12 precursor in aqueous solution [29]. The closo-dodecaboric acid enabled the synthesis of a range of metal–B12 H12 compounds through reactions between the acid and a base, oxide, carbonate or free metal. The resulting metal–B12 H12 was almost always formed as a strongly hydrated complex due to its synthesis in aqueous solution. Other closo-borane precursors may be used to make the Bn Hn –acid and resulting metal–Bn Hn (n = 6–12).

Dodecahydroxy-closo-dodecaborate, B12 (OH)2− 12 , is an versatile compound to which independent chains can be attached, enabling the synthesis of a wide selection of derivatives [52]. Cs2 B12 (OH)12 was prepared by heating a suspension of Cs2 B12 H12 in 30% hydrogen peroxide at reflux temperature for 3 days. Here, a precipitate is collected by refrigeration, then filtered and dried to yield a pure white powder of Cs2 B12 (OH)12 [52]. Partial halogenation or hydroxylation is also possible when the exchange reaction is not taken to completion, i.e., B12 H12−x Fx2− [51] or B12 H12−x OHx2− [53]. Significant efforts in the synthesis of boranes and carboranes for many different purposes have been conducted over the years. 2− 2− Table 1 lists a number of known B10 X10 and B12 X12 species, arranged by cation. These anions are the most common of the closoborane family, as such Bn Hn (n = 6–9, 11) are not included in this list, but can be found elsewhere [54].

2.5. Route V: Carborane synthesis

3. Structure and coordination in solid metal boranes

The previously mentioned routes all focus on the synthesis of metal closo-boranes, where the metal can vary, depending on the route. Closo-carboranes, the electronically neutral icosahedra, are 2− by two CH+ verobtained by replacing two BH vertices of B12 H12 tices. The C-atoms act as reaction centers and allow vast numbers of organic substitution reactions. There are several types of closocarboranes, i.e., C2 Bn Hn+2 (n = 4, 5, 8, 10), which are often found as three isomers: ortho, meta and para, based on carbon position [34,35]. Typically ortho-carborane is prepared by straightforward and benign conditions, while the meta and para isomers require higher temperatures. In the present work, only a short description of the synthesis for the twelve-vertex closo-C2 B10 H12 and its isomers is provided. Synthesis of the ortho-(1,2)carborane, C2 B10 H12 , is performed by the reaction of acetylene with decaborane [36]. Thermolysis of ortho-C2 B10 H12 at 465–500 ◦ C produces the metaisomer in high yield, while higher temperatures give an equilibrium mixture of para-(1,12)C2 B10 H12 and meta-(1,7)C2 B10 H12 [37]. Separation of the para and meta isomers can be performed by column chromatography, yielding 99.5% para-C2 B10 H12 [38]. Note that the addition of carbon changes the charge of the closo-borane anion. Thus, the charge can be altered from, e.g., [B12 H12 ]2− to [CB11 H12 ]− and [C2 B10 H12 ].

Many higher metal boranes are synthesized in water, and as such, many are known to complex with water. Typically, water first coordinates to the positively charged metal cation, and further water can usually also be absorbed into the crystal lattice, not coordinated directly to the cation. Deliquescence is also noted to occur, where powdered compounds quickly liquefy by absorption of water from the air [22]. In these compounds, water can exist in a large variety of coordination environments (Fig. 3). Typically, divalent cations will coordinate to 6 water molecules (e.g., Ba2+ , Cd2+ , Co2+ , Cu2+ , Fe2+ , Mg2+ , Mn2+ , Ni2+ , Zn2+ ) [46,51,53,59,61]; however, this is not always the case, i.e., Ca2+ may coordinate 7, and Sr2+ 8, water molecules [46,51]. Monovalent cations will likely coordinate to less water molecules, often forming bridged bonds to water molecules between two cations (Fig. 4). Similar trends also exist when acetonitrile is coordinated, where divalent Co2+ , Ni2+ and Pd2+ form 6-coordinate (octahedral) complexes between CH3 CN and the cation [51]. Similar to water, ammonia can be coordinated to the metal cation of MB12 H12 or MB10 H10 to form ammine complexes, i.e., Li2 B12 H12 ·7NH3 , Rb2 [B10 H10 ]·5NH3 or Cs[Na(NH3 )6 ][B10 H10 ]·NH3 [68,75,86]. These compounds are potential candidates for solid-state ammonia storage, which are currently highly relevant because NH3 is considered as a future energy carrier, especially in Japan [87]. 2− Halogenated closo-dodecaborates (B12 X12 , X = F, Cl, Br, I) have recently been highlighted as promising weakly coordinating anions (WCA) [27]. A weakly coordinating anion, sometimes classed as “non-coordinating”, will typically show weaker coordination to a cation than a solvent [88]. Ideal WCAs have a high degree of charge delocalisation over the entire anion, low charge and a large size, properties which are all typical of closo-dodecaborates [73]. This property enables the rich coordination chemistry in metal–B12 X12 compounds between cations and guest molecules, but can prove very problematic if a solvent-free compound is sought. Recently, the synthesis of solvent-free MgB12 H12 has been a focus of numerous studies [25,32,89,90]. Magnesium has a relatively high charge density (Fig. 4) and coordinates strongly to water or other solvent molecules (i.e., CH3 OH). As such, thermal treatment not only results in the desolvation of solvents, but also adverse decomposition reactions, releasing hydrogen gas [90]. It is believed that dehydrogenation is greatly facilitated by the formation of dihydrogen bonds between hydrogen in the solvent, H␦+ , 2− and hydrogen in the B12 H12 anion, Hı− [90]. A similar trend was observed in ammine metal borohydrides between Hı+ in NH3 and H␦− in BH4− , i.e., Mn(BH4 )2 -nNH3 [91]. Although this issue has been thoroughly investigated for magnesium dodecaborane it is likely to be problematic for other cations with comparable, or higher, charge densities.

2.6. Route VI: Functionalization While the closo-borane cation can be tailored using Routes I–IV, there are numerous ways to alter and change the anion to make functionalized boron cages. Substituted borane cages offer the ability to tune the properties of the material and have potential applications in many areas, such as weakly coordinating anions (WCA) [27,39], boron neutron capture therapy or drug delivery systems [40,41], as components for radio imaging [42] or ion conductivity [43,44]. Exchanging the hydrogens with halogens (X = F− , Cl− , Br− , I− ) changes the size of the boron-anion and thus changes the charge density and charge delocalization, which is correlated to the coordination ability. Halogenation of metal closo-boranes can be achieved by reactions with elemental halogens [45–47]. Chlorination is performed by passing chlorine-gas through aqueous or alcoholic 2− 2− solutions of B12 H12 or B10 H10 precursors, while bromination is achieved by mixing an aqueous solution with liquid Br2 . Iodination can be performed by reaction of excess I2 and ICl in methylenechloride, CH2 Cl2 , at 180 ◦ C [48]. Flourination was demonstrated by stirring K2 B12 H12 in liquid anhydrous HF at 70 ◦ C for 14 h, a pro2− to B12 H8 F42− in high yield [49]. cedure known to convert B12 H12 The mixture was then treated with 20% F2 /N2 for 72 h at 25 ◦ C to fully fluorinate the boron cage [50,51].

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Table 1 List of metal B10 H10 2– and B12 H12 2– compounds and halide derivatives arranged by cation. Metal

Compound

Ref

Ag

Ag(CH3 CN)2 B12 H12 Ag2 (C7 H8 )6 B12 F12 Ag2 (CH2 Cl2 )4 B12 F12 Ag2 (CH3 CN)4 B12 F12 Ag2 (CH3 CN)5 B12 F12 Ag2 (CH3 CN)8 B12 F12 Ag2 (H2 O)4 B12 F12 Ag2 B10 H10 Ag2 B10 H10 ·DMF Ag2 B12 Br12 Ag2 B12 Cl12 Ag2 B12 F12 Ag2 B12 H12 (Al(H2 O)6 )2 (B12 H12 )3 .16H2 O (Al(H2 O)6 )2 (SO4 )(B12 H12 )2 ·15H2 O (Et2 Al)2 B12 Cl12 (Me2 Al)2 B12 Cl12 Al6 (OH)10 (H2 O)14 (B12 H12 )4 ·28H2 O Ba(H2 O)5 B12 F12 Ba(H2 O)6 B12 H12 BaB12 H12 Bi(H2 O)2 (OH)B12 H12 Ca(H2 O)7 B12 F12 Ca(H2 O)7 B12 H12 ·H2 O CaB12 H12 Cd(H2 O)6 (H3 O)2 (B12 H12 )2 ·6H2 O Cd(H2 O)6 B12 H12 (Co(NH3 )6 )2 (B12 H12 )3 .6H2 O Co(CH3 CN)6 B12 F12 Co(H2 O)6 (H3 O)2 (B12 H12 )2 ·6H2 O Co(H2 O)6 B10 H10 Co(H2 O)6 B10 H10 ·2H2 O Co(H2 O)6 B12 F12 Co(H2 O)6 B12 H12 ·1.6H2 O Co(H2 O)6 B12 H12 ·6H2 O CoB10 H10 .0.06H2 O (Cr(H2 O)6 )2 (B12 H12 )3 ·15H2 O (Cr(NH3 )6 )2 (B12 H12 )3 ·7H2 O Cr(H2 O)6 (H5 O2 )(B12 H12 )2 ·6H2 O Cs(N(CH3 )4 )B12 H12 ·H2 O Cs2 (CH3 CN)B12 F12 Cs2 (H2 O)B12 F12 Cs2 (HF)B12 F12 Cs2 B10 H10 Cs2 B12 Br12 Cs2 B12 Br12 ·2CH3 CN Cs2 B12 Br12 ·2H2 O Cs2 B12 Br12 ·H2 O Cs2 B12 Cl10 H2 Cs2 B12 Cl12 Cs2 B12 Cl12 ·2CH3 CN Cs2 B12 Cl12 ·SO2 Cs2 B12 F12 Cs2 B12 H12 Cs2 B12 H12 Cs2 B12 I12 Cs2 B12 I12 ·2CH3 CN Cs3 (AsF6 )(B12 F12 ) Cs3 BF4 B12 H12 Cs3 BH4 B12 H12 Cs3 BrB12 H12 Cs3 ClB12 H12 Cs3 IB12 H12 Cs3 NO3 B12 H12 CsCuB10 H10 Cs3 La(H2 O)8 (B12 H12 )3 ·11H2 O CsNa(NH3 )6 B10 H10 ·NH3 Cu(H2 O)5.5 B12 H12 ·2.5H2 O Cu(H2 O)6 (H3 O)2 (B12 H12 )2 ·6H2 O Cu(H2 O)6 B12 H12 ·4H2 O Cu2 B10 H10 Cu2 B12 H12 Eu2 (B12 H12 )3 ·11H2 O EuB10 H10 EuB12 H12

[55] [51] [51] [51] [51] [51] [51] [29] [56] [57] [51,57] [51] [29] [29] [53] [58] [58] [59] [51] [46] [29,60] [53] [51] [46] [32] [53] [46] [29] [51] [53] [29] [61] [51] [29] [46] [29] [53] [29] [53] [46] [51] [51] [51] [29] [46,62,63] [46] [46] [46] [46] [46,57,63] [46] [64] [51] [57] [46] [46,63] [46] [65] [46] [46] [46] [46] [46] [46] [66,67] [59] [68] [46] [53] [59] [69] [29] [29] [70] [70]

Al

Ba

Bi Ca

Cd Co

Cr

Cs

Cu

Eu

5

Table 1 (Continued) Metal

Compound

Ref

Fe

Fe(H2 O)6 (H3 O)2 (B12 H12 )2 ·6H2 O Fe(H2 O)6 B10 H10 ·2H2 O Fe(H2 O)6 B12 H12 (Et2 Ga(OEt2 )2 )2 B12 Cl12 HgB10 H10 HgB12 H12 (Ho(H2 O)9 )2 (B12 H12 )3 ·15H2 O Ho(H2 O)9 (H3 O)3 (B12 H12 )3 ·9H2 O (Et2 In)2 B12 Cl12 (In(H2 O)6 )2 (B12 H12 )3 ·15H2 O K2 (CH3 CN)2 B12 F12 K2 (H2 O)2 B12 F12 K2 (H2 O)4 B12 F12 K2 (H2 O2 )1.5 (H2 O)0.5 B12 F12 K2 (HF)3 B12 F12 K2 B10 H10 K2 B12 Br12 ·3H2 O K2 B12 Cl12 ·8SO2 K2 B12 F12 K2 B12 H12 K3 (AsF6 )(B12 F12 ) K3 BrB12 H12 K3 IB12 H12 (La(H2 O)10 )2 (B12 H12 )3 ·15H2 O (La(H2 O)9 )(H3 O)Cl2 B12 H12 ·H2 O (La(H2 O)9 )2 (B12 H12 )3 ·7H2 O Li2 (H2 O)4 B12 F12 Li2 (H2 O)7 B12 H12 Li2 (SO2 )8 B12 Cl12 Li2 (SO2 )8 B12 Cl12 Li2 B10 H10 ·H2 O Li2 B12 F12 Li2 B12 H12 Li2 B12 H12 ·7NH3 LiCB11 H12 LiK(H2 O)4 B12 F12 LiNaB12 H12 (Li0.7 Na0.3 )3 BH4 B12 H12 Mg(H2 O)6 B12 F12 Mg(H2 O)6 B12 H12 ·6H2 O MgB12 H12 Mg(H2 O)6 (CB11 H12 )2 Mg(THF)6 (CB11 H12 )2 Mg(Monoglyme)3 (CB11 H12 )2 Mg(Diglyme)3 (CB11 H12 )2 Mn(H2 O)6 (H3 O)2 (B12 H12 )2 ·6H2 O Mn(H2 O)6 B10 H10 ·2H2 O Mn(H2 O)6 B12 H12 Mn(H2 O)6 B12 H12 ·6H2 O MnB12 H12 ·0.07H2 O Na2 (H2 O)4 B12 F12 Na2 (H2 O)4 B12 H12 Na2 (H2 O)8 B12 Br12 Na2 B10 H10 Na2 B10 H10 ·2H2 O Na2 B12 Br12 Na2 B12 Cl12 Na2 B12 Cl12 ·4SO2 Na2 B12 H12 Na2 B12 I12 Na3 BH4 B12 H12 NaCB11 H12 (NH4 )2 B12 Br12 ·3H2 O (NH4 )2 B12 H12 (NH4 )3 BrB12 H12 (NH4 )3 IB12 H12 (NH4 ·H2 O)2 B12 F12 Ni(C2 H4 (NH2 )2 )3 B12 H12 Ni(CH3 CN)6 B12 F12 Ni(H2 O)6 (H3 O)2 (B12 H12 )2 ·6H2 O Ni(H2 O)6 B10 H10 ·2H2 O Ni(H2 O)6 B10 H10 ·3H2 O Ni(H2 O)6 B12 F12 Ni(H2 O)6 B12 H12 ·0.25H2 O Ni(H2 O)6 B12 H12 ·6H2 O Ni(NH3 )6 B12 H12 ·0.5H2 O NiB12 Br12 .7CH3 CN

[53] [61] [53] [58] [71] [53] [53] [53] [58] [53] [51] [51] [51] [51] [51] [72] [46] [64] [51] [31,46] [65] [46] [46] [46] [46] [53] [51] [46] [73] [64] [29] [50] [31,74] [75] [76] [51] [77] [78] [51] [46] [32] [79] [79] [79] [79] [53] [61] [29] [46] [29] [51] [46,59] [46] [72] [29] [57] [57] [64] [31,57] [80] [78] [76] [46] [46] [46] [46] [51] [81] [51] [53] [61] [29] [51] [29] [53] [29] [82]

Ga Hg Ho In K

La

Li

Mg

Mn

Na

NH4

Ni

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6 Table 1 (Continued) Metal

Pb Pd Pr Rb

Sc Si

Sr

Tl U Yb Zn

4.1. Ion conductivity Compound

Ref

NiB12 Cl12 .6CH3 CN NiB12 H12 ·0.8H2 O Pb(H2 O)3 B12 H12 ·3H2 O Pd(CH3 CN)6 B12 F12 ·2CH3 CN (Pr(H2 O)9 )2 (B12 H12 )3 ·15H2 O Rb2 (H2 O)4 B12 F12 Rb2 (H2 O)B12 F12 Rb2 B10 H10 Rb2 B10 H10 ·1.5H2 O Rb2 B12 Br12 ·H2 O Rb2 B12 Cl12 ·4SO2 Rb2 B12 H12 Rb3 BrB12 H12 Rb3 ClB12 H12 Rb3 IB12 H12 Sc(H2 O)7 (HCO3 )B12 H12 ·6H2 O (Et3 Si)2 B12 Br12 ·DCB (Et3 Si)2 B12 Cl2 (Et3 Si-H-SiEt3 )2 B12 Br2 (Et3 Si-H-SiEt3 )2 B12 Cl12 (Et3 SiN2 (CH3 )4 )Et3 SiB12 Br12 Sr(H2 O)8 B12 H12 SrB12 H12 SrB12 H12 ·7H2 O Tl2 B10 H10 Tl2 B12 H12 UO2 (H2 O)5 B12 H12 ·6H2 O UO2 B10 H10 ·7CO(NH2 ) YbB10 H10 YbB12 H12 Zn(H2 O)2 (CH3 CN)4 B12 F12 ·H2 O Zn(H2 O)6 (H3 O)2 (B12 H12 )2 ·6H2 O Zn(H2 O)6 B10 H10 ·2H2 O Zn(H2 O)6 B12 H12 ·6H2 O

[82] [29] [53] [51] [53] [51] [51] [72] [83] [46] [64] [46] [46] [46] [46] [59] [57] [57] [57] [57] [57] [46] [60] [84] [29] [29,53] [85] [85] [70] [70] [51] [53] [61] [46]

The charge density of a particular cation is a function of a number of key parameters, including its charge, ionic radius and coordination number (Fig. 4) [92]. The charge density of a cation has a direct correlation to its ability to coordinate ligands such as water. Thus, it is more difficult to remove coordinated solvents from compounds with a high charge density cation. The coordination number (CN) of the cation has a significant impact on its charge density and lower CNs lead to higher charge densities [92]. Whilst many factors other than the charge density of the metal may impact the ability for a solvent to coordinate, Fig. 4 does help illustrate why elements with high charge density, e.g., Mg2+ and Al3+ , typically decompose before desolvation and elements with low charge density, e.g., Cs+ , Ag+ and K+ , can be desolvated without decomposing, resulting in a pure, waterfree metal borane (See Table 1), i.e., MgB12 H12 ·4H2 O decomposes before it is able to be fully dehydrated, whereas Li2 B12 H12 ·4H2 O is able to be fully dehydrated without decomposition at T = 245 ◦ C.

4. Properties and applications of metal boranes Some of the first practical applications for boranes were for militaristic purposes as either rocket or aircraft fuels. Ultimately the higher boranes were not well suited for warfare applications, and since the end of the cold war, research has been aimed at peaceful applications, e.g., medicine, solid-state electrolytes, boride precursors and reducing agents in organic chemistry. These applications are still growing as metal boranes are more thoroughly characterized. For instance, Li2 B12 H12 has recently been found to exhibit photophysical properties that indicate a potential application as a luminescent dye in transparent head-up displays [93].

Recent research into higher metal boranes has focused on their possible use as solid state ion conductors. Polymorphic transitions in these compounds occur at specific transition temperatures, resulting in rapid reorientational dynamics of the borane anions. Fast cationic conductivity is often observed in the high temperature polymorphs [94]. An overview of the structural transition temperatures for a range of metal–boron compounds is shown in Fig. 5. Rapid anion dynamics are well-known in other types of compounds based on PO43− , SO42− , or BF4− anions. Often these reorientations are coupled by fast cation conduction, but the exact mechanism behind this can vary. Typical models involve either a ‘paddle-wheel’ mechanism and/or a percolation-type mechanism [95]. Some discussion has surrounded the definition of the ‘paddle-wheel’ mechanism [96], which may be considered a correlation between anion dynamics and cation jump frequency. The mechanism describes the cation migration between sites and suggests that energy barriers for migration could be lowered by anion rotation due to anion shape, displacement and localized cationanion attractions. 1 H NMR measurements of the high temperature polymorphs of Li2 B12 H12 and Na2 B12 H12 show an increase by approximately two orders of magnitude in the reorientational jump rate [97]. Impedance measurements of Na2 B12 H12 reveal Na+ conductivity above its order–disorder phase transformation (256 ◦ C) rivaling that of current solid state, ceramic based, Na-battery electrolytes (order of 0.1 S cm−1 ) [98]. The properties of Na2 B12 H12 can also be altered by cation and anion modifications [76–78]. The BH4 − anion is known to undergo rapid reorientation in the high-temperature (HT) polymorphs of MBH4 (M = Li, Na, K, Rb, Cs) [102]. In the case of borohydrides, there is a clear trend, where larger cations have lower transition temperatures. The MB12 H12 (M = Li, Na, K, Rb, Cs) compounds follow a more complex trend. The transition temperature is related to cation size, anion size and structural coordination. Solid solutions of Lix Na2−x B12 H12 have also been investigated and shown to have transition temperatures in between or below their constituents’ temperatures [43,77]. The transition temperatures of Li2 B12 H12 , Na2 B12 H12 and their solid solutions have all been determined with DSC under different heating rates, which could result in the deviations between LiNaB12 H12 (Ref. [77]) and Lix Na2−x B12 H12 (Ref. [43]) in Fig. 5. It is also notable that Na2 B10 H10 has a transition temperature in between those of NaBH4 and Na2 B12 H12 , demonstrating that the transition temperature may depend on anion size, where larger anions undergo reorientation at higher temperatures. This is also true for Na2 B12 X12 (X = H, Cl, I) [43]. However, this trend does not hold for the case of Cs2 B12 X12 (X = H, Cl, Br, I) [45] so other factors such as localized cation–anion attractions may also impact the transition temperature. Recently, it was demonstrated that the phase transition temperature can be tuned by halogenation of the boron cage and by cation substitution, i.e., mixed-cation closo-boranes with a phase transition in the temperature range between the two untreated compounds [43,77]. Although the larger halogenated anions give rise to phase transitions at higher temperatures, it does demonstrate the ability to tune the anion/cation size ratio, which is linked to the HT phase transition. Na2 B10 H10 , which transforms to a HT polymorph at 90 ◦ C, shows superionic conductivity ( = 0.01 S cm−1 ), superior to many known complex hydrides [100]. The high conductivity and favorable properties of the Na2 B10 H10 may encourage applications of Na-ion battery technologies. Although the phase transformation for Na2 B10 H10 (90 ◦ C) is much lower than Na2 B12 H12 (256 ◦ C), it does not indicate a systematic trend, as Li2 B10 H10 and Li2 B12 H12 have similar transition temperatures (367 and 355 ◦ C, respectively), see Fig. 5 [97,101]. Exceptional ion conductivity has also been recently

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7

Fig. 3. Various coordination environments of metal–water complexes observed in borane crystals. Water (oxygen: red) metal cations (blue) [46,51,65,85].

shown in Li and NaCB11 H12 [76]. This is another example of how transition temperatures can be tuned by different anions. Here, LiCB11 H12 and NaCB11 H12 , have transition temperatures of 127 and 107 ◦ C, respectively and demonstrate some of the best Li and Na ion conductivity results obtained (See Fig. 6). Mg(CB11 H12 )2 has also been proposed as an electrolyte for Mg-batteries [79]. It has been investigated as a tri- and tetraglyme adduct and shows remarkable properties: non-corrosive towards the cell, high anodic stability and performance. These promising results should lead to further studies with CB11 H12 -based compounds as solid-state electrolytes. Another method of shifting high ion conductivity towards RT is to avoid the high temperature transition entirely. Whilst the pure 2− 2− and B12 H12 only show high ionic mobility above the metal B10 H10 order-disorder phase transition, the mixed-anion Na3 BH4 B12 H12 and (Li0.7 Na0.3 )3 BH4 B12 H12 conduct Na+ and/or Li+ at high rates already from RT (10−3 S cm−1 ), approaching that of Na-␤-alumina

[78]. An explanation for the improvement may be that Na-rich regions, within the crystal structure, layered with Na-poor regions mitigate a 2D conduction pathway. In Fig. 6 the ion conductivity of the most promising Li and Na materials is presented. The recent progress in mixed-anion and mixed-cation metal boranes demonstrates the capability to tune ion conductivity in these solid state materials. 4.2. Boranes in medicine Another purpose for boranes and boron clusters is found in boron neutron capture therapy (BNCT) for cancer treatment due to the large neutron scatter cross section of 10 B and also since their chemical stability in biological systems make them relatively nontoxic. Here, disodium mercapto-undecahydrocloso-dodecaborate, Na2 B12 H11 SH, and NaB12 H11 NH3 , relatives

Fig. 4. Charge density of selected cations typically when in coordination number 6 [92].

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Fig. 5. Reorientational motion transition temperatures of selected metal–boron compounds [43,45,76,77,97,99–101].

of Na2 B12 H12 , show promise as boron-delivery systems for this minimally invasive cancer treatment [106]. The high stability, water solubility, low toxicity and high boron content make Na2 B12 H12 and Na2 B10 H10 ideal candidates in BNCT. They, and their derivatives have been studied extensively [107–110]. However, combining these beneficial properties with tumor selectivity, the ability to pass the blood-brain barrier (BBB) and their cost requires further research [41,111]. In fact, the charged borane anions may negatively influence the transport, especially through lipophilic barriers, such as cell membranes and the BBB [40]. Closo-carboranes have a lower boron-content, but allow for vast functionalization through the introduced carbon reaction centers. This opens numerous possibilities to synthesize boron containing compounds, e.g., amino acids, nucleosides, porphyrins etc., which could all be applied to BNCT and other medicinal purposes [111,112]. Furthermore, the lower charge of carboranes, i.e., −1 in CB11 H12 1− and 0 in o-, p-, m-C2 B10 H12 , provides a more lipophilic component, which enhances the ability to cross the BBB [40]. Nevertheless Na2 B12 H11 SH, commonly denoted BSH, and 4dihydroxyborylphenylalanine, denoted BPA, are so far two of the most successful compounds in clinical trials [113]. The first clinical trial with BSH was performed in 1968 by Hatanaka and Nakagawa in Japan [114,115] and BPA-fructose (BPA-F) was used for BNCT in 1994 [116]. Both compounds were synthesized over 50 years ago, and one of the great needs in BNCT research is the development of new and more selective boron delivery agents. The ideal boron carrying compound should be non-toxic and very selective in targeting tumor cells. Thus, it should not accumulate in normal tissue included in the radiation field. This ideal compound should also penetrate the BBB to reach the microscopic extensions of tumor and it should clear from blood and normal brain tissue before BNCT radiation is delivered [117]. BPA-F and BSH satisfy some of the requirements of the ideal boron carrier, but currently no compound fulfills all the requirements. However, there are promising candidates under development, where boron or polyhedral boron agents are incorporated into tumor targeting molecules including peptides, sugars, proteins, liposomes, nucleosides and porphyrins [118–122]. A recent review of BNCT underlines that the success of

Fig. 6. Li+ and Na+ ion conductivity in selected well-known materials [76–78,98,100,103–105].

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BNCT is not solely linked to the boron carrying drug, but also: the pharmaceutical industry having to be involved in drug development, which has so far it has been reluctant to do so. Furthermore, large scale trials are necessary, which is only feasible with the presence of a significant number of neutron facilities in specialized hospital-based institutions [113]. If some of these requirements are met there may be a bright future for BNCT. The wealth of new boranes being investigated for alternate applications may also hold promise for BNCT, especially the carboranes. 4.3. Carborane and borane-based polymers Borane-based polymers encompass a diverse field, which started during early investigations of decaborane where it was found that it will react with Lewis bases, L, to form a diadduct B10 H12 L2 [123]. When bidentate ligands are used it is possible to form polymeric chains. This synthesis method is similar to the one used to form carboranes and other boranes as described earlier. The polymers from decaborane with bidentate ligands are not polymers of closed icosahedral boron cages, but rather nido-shaped decaborane with ligands in between. However, icosahedral closo-carboranes and -boranes can form polymers, where the closed boron cages are linked via bridging hydrogen. In the case of closo-boranes, polymerization must go through the superacidic (H3 O)2 B12 H12 oily solution, which polymerizes under thermal treatment by Reaction (6) [124], whereas other derivatives, primarily carboranes, readily form polymers, when C-atoms act as linkers between cages. 3[B12 H12 ]2– + 2H+ → [B36 H34 ]4– + 2H2

(6)

It is noted that the polymerization of boranes and especially carboranes is somehow similar to polymerization of fullerenes [13,14], likely owing to the spherical shapes of polyhedral boranes and fullerenes. Borane and carborane containing polymers have so far been most widely employed as precursors for boron carbide, boron nitride and related ceramic materials [125]. However, polymers based on borane and carborane may be used as building blocks for nanosized porous ceramics that can be used for gas separation or storage and catalyst support (similar to the properties of metal organic frameworks, MOF’s). The self-assembly of carborane-based polymers can produce porous ceramics with different morphologies depending on the solvent [126]. The ability to design nanostructures with carborane polymers may also play a role in BNCT, where the size of reagents is critical for selective delivery to the tumor cells [127]. On a similar note, organic derivatives of closo-boranes and carboranes may also be used for liquid crystal materials, where the icosahedral clusters serve as effective structural elements in rodlike molecular chains. Combinations of certain building blocks can vary the properties of the liquid crystals, allowing customization of the UV transparency or refractive indices [128]. Self-assembled monolayers have been demonstrated with decaborane- and carboranethiols on coinage metal surfaces, which may have widespread applications as inorganic building blocks in surface chemistry, although research is still at a fundamental level [129–131]. 4.4. Higher boranes in hydrogen storage Significant research on metal borohydrides has been undertaken since the turn of the millennium, mainly because of the very high volumetric and gravimetric hydrogen densities in this class of materials. This makes them ideal candidates for solid state storage of hydrogen/renewable energy. However, most borohydrides have complex stepwise decomposition pathways typically accompanied 2− 2− and B12 H12 , which can by formation of anions such as B3 H8− , B10 H10 hamper reversibility. Generally the alkali and alkaline earth metal

9

borohydrides have dominantly ionic metal–B bonds, while transition metal borohydrides are more covalent. This relation together with the Pauling electronegativity (or ionic potential [132]) of the metal is linked to the stability and decomposition temperature of the borohydride [133]. The released gas depends on these factors [134], i.e., unstable/covalent metal borohydrides (xP > 1.5) usually release a mix of diborane (B2 H6 ) and hydrogen, whereas more stable/ionic metal borohydrides (xP < 1.5) only release hydrogen. However MB12 H12 decomposition intermediates can be formed in both cases [134]. In fact, the formation of MB12 H12 intermediates during decomposition may be described as a reaction between MBH4 and B2 H6 , similar to Route I in Fig. 2 [135]. MB12 H12 has been observed during decomposition of numerous borohydrides, including LiBH4 [24,74], NaBH4 [136] and Ca(BH4 )2 [137]. Decomposition products also depend heavily on desorption conditions, i.e., temperature, hydrogen backpressure and particle size of the sample [138]. For example, a LiBH4 –MgH2 –Al (4:1:1) composite contains either 19 or 9 mol% Li2 B12 H12 after three hydrogen release cycles at p(H2 ) = 0.15 or 5.0 bar, respectively [139]. This has also been demonstrated in studies of Ca(BH4 )2 decomposition by varying hydrogen back pressure [140]. MgH2 or Ti and Nb based additives have also shown to suppress the formation of CaB12 H12 in Ca(BH4 )2 , which improves reversible hydrogen storage capacity [141–143]. The role of transition metal additives has also been highlighted in LiBH4 –MgH2 to improve reaction times [144]. Whilst B3 H8 − compounds may enable rehydrogenation in some cases 2− [145], the B12 H12 anion has been regarded to inhibit rehydrogenation in metal borohydrides [139,146–148]. A novel strategy is to only partially decompose a metal borohydride by prolonged heating at a lower fixed temperature relative to what would provide release of the full hydrogen content. This can be illustrated by magnesium borohydride, Mg(BH4 )2 , which can undergo multiple polymorphic transitions during heating [149,150] under pressure [151] and it can even melt [152]. The decomposition mechanism of magnesium borohydride is quite complex and highly dependent on reaction conditions [149,153–158]. Amorphous MgB12 H12 , was reported during the decomposition of Mg(BH4 )2 in vacuum [153], but Mg(B3 H8 )2 was found as the major decomposition product after 5 weeks at 200 ◦ C in vacuum [159]. In another study, Mg(BH4 )2 decomposed to Mg(B3 H8 )2 at 265 ◦ C, but above T > 285 ◦ C none of the boranes Mg(BH4 )2 , Mg(B3 H8 )2 or MgB12 H12 were observed [160]. The temperature used to decompose a borohydride influences its decomposition products, which is the same behavior observed during synthesis of higher boranes (Fig. 2), where higher reaction temperatures result in more stable borane products. The hydrogen release from MB12 H12 compounds is slow and often restricted to high temperatures, which is unfavorable for hydrogen storage applications. For instance, Li2 B12 H12 decomposes over a large temperature range from 250 ◦ C to above 500 ◦ C, forming an amorphous hydrogen-poor Li2 B12 H12-x composition [22]. It is common that extreme reaction conditions are required to rehydrogenate decomposed metal borohydrides, i.e., 400 bar, 270 ◦ C, 48 h for Mg(BH4 )2 [158]. Production of metal borohydrides from the expected decomposition products has also been demonstrated using 700–950 bar H2 and elevated temperature, although reaction time is relatively long (∼100 h) even at these conditions [161,162]. Naturally, the aim of much research has been to limit the formation of MB12 H12 , in order to retain reversible, lower-temperature hydrogen storage capacity. Recent results are promising, for example, a mixture of Mg(B3 H8 )2 ·2THF–MgH2 (1:2) was heated to 200 ◦ C under p(H2 ) = 50 bar for 2 h and the product indicated a quantitative conversion to Mg(BH4 )2 [163]. Many problems in understanding the decomposition of metal borohydrides have come from the fact that the higher metal borane reaction products are difficult to characterize and identify, especially because they are typically non-crystalline [164]. This makes their role in the

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hydrogen release and uptake pathways difficult to elucidate. A recent study even indicates that the thermal decomposition of pure alkali metal dodecaborates is unlike the decomposition of dodecaborates formed from metal borohydrides [165]. Further research into metal boranes will assist with understanding their role in the decomposition of metal borohydrides, possibly leading to lower hydrogen release temperatures and better hydrogen reversibility. 5. Conclusions The chemistry of metal boranes is diverse and branches into many related, yet very different, research fields. A broad scope of applications calls for new materials, which pushes research progress to develop new synthesis methods. Herein we present and discuss detailed synthesis routes as well as the rich coordination chemistry that most metal boranes display. It is fascinating that many of the current synthesis methods and applications were developed over 50 years ago, yet fundamental questions regarding structure, dynamics and reaction mechanisms continue to intrigue scientists today in many research fields. Generally, significant developments have been made, and herein we describe some of the recent progress in the science and application of metal boranes of medicine, polymers, ion conductivity and energy storage. Closo-boranes and especially carboranes have shown promise in the fields of BNCT, polymers, gold surface chemistry and ion conductivity with further research likely to lead to fruitful progress. Closo-boranes have also shown promise in hydrogen storage and yield other interesting properties, such as photoluminescence or as liquid crystal materials, which open new fields of exploration in future research. Acknowledgments The work was supported by the Danish National Research Foundation, Center for Materials Crystallography (DNRF93), The Innovation Fund Denmark (project HyFill-Fast), the Danish Research Council for Nature and Universe (Danscatt) and the Danish Council for Independent Research (DFF Mobility 1325-00072). We are grateful to the Carlsberg Foundation. References [1] I.B. Sivaev, V.I. Bregadze, S. Sjöberg, Collect. Czechoslov. Chem. Commun. 67 (2002) 679–727. [2] A. Stock, Hydrides of Boron and Silicon, Cornell University Press, Ithaca, NY, 1933. [3] A.B. Burg, H.I. Schlesinger, J. Am. Chem. Soc. 59 (1937) 780–787. [4] H.I. Schlesinger, H.C. Brown, J. Am. Chem. Soc. 62 (1940) 3429–3435. [5] W.H. Eberhardt, B. Crawford Jr., W.N. Lipscomb, J. Chem. Phys. 22 (1954) 989–1001. [6] W.M. Lipscomb, Boron Hydrides, W.A. Benjamin Inc., New York, 1963. [7] E.L. Muetterties, The Chemistry of Boron and Its Compounds, First Ed., John Wiley and Sons. Benjamin Inc., New York, 1967. [8] D.M. Schubert, Kirk-Othmer Encyclopedia of Chemical Technology, John Wiley & Sons, Inc., New York, 2000. [9] J.D. Clark, Ignition!: An Informal History of Liquid Rocket Propellants, Rutgers University Press, New Brunswick, NJ, 1972. [10] J.J. Rockwell, A. Herzog, T. Peymann, C.B. Knobler, F.M. Hawthorne, Curr. Sci. 78 (2000) 405–409. [11] R. Tycko, Solid State Nucl. Magn. Reson. 3 (1994) 303–314. [12] L. Maidich, D. Pontiroli, M. Gaboardi, S. Lenti, G. Magnani, G. Riva, P. Carretta, C. Milanese, A. Marini, M. Riccò, S. Sanna, Carbon 96 (2016) 276–284. [13] M. Riccò, M. Belli, M. Mazzani, D. Pontiroli, D. Quintavalle, A. Jánossy, G. Csányi, Phys. Rev. Lett. 102 (2009) 145901. [14] D. Pontiroli, M. Aramini, M. Gaboardi, M. Mazzani, A. Gorreri, M. Riccò, I. Margiolaki, D. Sheptyakov, Carbon 51 (2013) 143–147. [15] J.L. Boone, J. Am. Chem. Soc. 86 (1964) 5036. [16] F. Klanberg, D.R. Eaton, L.J. Guggenberger, E.L. Muetterties, Inorg. Chem. 6 (1967) 1271–1281. [17] F. Klanberg, E.L. Muetterties, Inorg. Chem. 5 (1966) 1955–1960. [18] A.R. Pitochelli, F.M. Hawthorne, J. Am. Chem. Soc. 82 (1960) 3228–3229. [19] M.F. Hawthorne, A.R. Pitochelli, J. Am. Chem. Soc. 81 (1959) 5519. [20] W.V. Hough, L.J. Edwards, A.D. McElroy, J. Am. Chem. Soc. 78 (1956) 689.

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