`color{green}(★)` The simplest boron hydride known, is diborane. It is prepared by treating boron trifluoride with `color{red}(LiAlH_4)` in diethyl ether.
`color{red}(4BF_3 + 3 LiAlH_4 → 2B_2H_6 + 3LiF + 3AlF_3)`
`color{green}(★)` A convenient laboratory method for the preparation of diborane involves the oxidation of sodium borohydride with iodine.
`color{red}(2NaBH_4 + I_2 → B_2H_6 + 2NaI + H_2)`
`color{green}(★)` Diborane is produced on an industrial scale by the reaction of `color{red}(BF_3)` with sodium hydride.
`color{red}(2BF_3 + 6NaH overset(450K)→ B_2H_6 + 6NaF)`
`color{green}(★)` Diborane is a colourless, highly toxic gas with a b.p. of `180 K`. Diborane catches fire spontaneously upon exposure to air. It burns in oxygen releasing an enormous amount of energy.
`color{red}(B_2H_6 + 3O_2 → B_2O_3 + 3H_2O ; \ \ \ \ Delta_cH^(⊖) = -1976 kJ mol^(-1))`
`color{green}(★)` Most of the higher boranes are also spontaneously flammable in air. Boranes are readily hydrolysed by water to give boric acid.
`color{red}(B_2H_6(g) + 6H_2O(l) → 2B(OH)_3(aq) + 6H_2(g))`
`color{green}(★)` Diborane undergoes cleavage reactions with Lewis bases `(L)` to give borane adducts, `color{red}(BH_3⋅L)`
`color{red}(B_2H_6 + 2N Me_3 → 2BH_3 NMe_3)`
`color{red}(B_2H_6 + 2 CO → 2BH_3⋅CO)`
`color{green}(★)` Reaction of ammonia with diborane gives initially `color{red}(B_2H_6.2NH_3)` which is formulated as `color{red}([BH_2(NH_3)_2]^(+) [BH_4]^(–))` ; further heating gives borazine, `color{red}(B_3N_3H_6)` known as “inorganic benzene” in view of its ring structure with alternate `color{red}(BH)` and `color{red}(NH)` groups.
`color{red}(3B_2H_6+6NH_3 → 3 [BH_2 (NH_3)_2]^(+) [BH_4]^(-) overset("Heat")→ 2B_3N_3H_6 + 12 H_2)`
`color{green}(★)` The structure of diborane is shown in Fig.11.2(a).
The four terminal hydrogen atoms and the two boron atoms lie in one plane. Above and below this plane, there are two bridging hydrogen atoms. The four terminal `color{red}(B-H)` bonds are regular two centre-two electron bonds while the two bridge `color{red}((B-H-B))` bonds are different and can be described in terms of three centre–two electron bonds shown in Fig.11.2 (b).
`color{green}(★)` Boron also forms a series of hydridoborates; the most important one is the tetrahedral `color{red}([BH_4]^(–))` ion. Tetrahydridoborates of several metals are known. Lithium and sodium tetrahydridoborates, also known as borohydrides, are prepared by the reaction of metal hydrides with `color{red}(B_2H_6)` in diethyl ether.
`color{red}(2MH + B_2H_6 → 2 M^(+) [BH_4]^(–) \ \ \ \ (M = Li \ \ \ \ "or" \ \ \ \ Na))`
`color{green}(★)` Both `color{red}(LiBH_4)` and `color{red}(NaBH_4)` are used as reducing agents in organic synthesis. They are useful starting materials for preparing other metal borohydrides.
`color{green}(★)` The simplest boron hydride known, is diborane. It is prepared by treating boron trifluoride with `color{red}(LiAlH_4)` in diethyl ether.
`color{red}(4BF_3 + 3 LiAlH_4 → 2B_2H_6 + 3LiF + 3AlF_3)`
`color{green}(★)` A convenient laboratory method for the preparation of diborane involves the oxidation of sodium borohydride with iodine.
`color{red}(2NaBH_4 + I_2 → B_2H_6 + 2NaI + H_2)`
`color{green}(★)` Diborane is produced on an industrial scale by the reaction of `color{red}(BF_3)` with sodium hydride.
`color{red}(2BF_3 + 6NaH overset(450K)→ B_2H_6 + 6NaF)`
`color{green}(★)` Diborane is a colourless, highly toxic gas with a b.p. of `180 K`. Diborane catches fire spontaneously upon exposure to air. It burns in oxygen releasing an enormous amount of energy.
`color{red}(B_2H_6 + 3O_2 → B_2O_3 + 3H_2O ; \ \ \ \ Delta_cH^(⊖) = -1976 kJ mol^(-1))`
`color{green}(★)` Most of the higher boranes are also spontaneously flammable in air. Boranes are readily hydrolysed by water to give boric acid.
`color{red}(B_2H_6(g) + 6H_2O(l) → 2B(OH)_3(aq) + 6H_2(g))`
`color{green}(★)` Diborane undergoes cleavage reactions with Lewis bases `(L)` to give borane adducts, `color{red}(BH_3⋅L)`
`color{red}(B_2H_6 + 2N Me_3 → 2BH_3 NMe_3)`
`color{red}(B_2H_6 + 2 CO → 2BH_3⋅CO)`
`color{green}(★)` Reaction of ammonia with diborane gives initially `color{red}(B_2H_6.2NH_3)` which is formulated as `color{red}([BH_2(NH_3)_2]^(+) [BH_4]^(–))` ; further heating gives borazine, `color{red}(B_3N_3H_6)` known as “inorganic benzene” in view of its ring structure with alternate `color{red}(BH)` and `color{red}(NH)` groups.
`color{red}(3B_2H_6+6NH_3 → 3 [BH_2 (NH_3)_2]^(+) [BH_4]^(-) overset("Heat")→ 2B_3N_3H_6 + 12 H_2)`
`color{green}(★)` The structure of diborane is shown in Fig.11.2(a).
The four terminal hydrogen atoms and the two boron atoms lie in one plane. Above and below this plane, there are two bridging hydrogen atoms. The four terminal `color{red}(B-H)` bonds are regular two centre-two electron bonds while the two bridge `color{red}((B-H-B))` bonds are different and can be described in terms of three centre–two electron bonds shown in Fig.11.2 (b).
`color{green}(★)` Boron also forms a series of hydridoborates; the most important one is the tetrahedral `color{red}([BH_4]^(–))` ion. Tetrahydridoborates of several metals are known. Lithium and sodium tetrahydridoborates, also known as borohydrides, are prepared by the reaction of metal hydrides with `color{red}(B_2H_6)` in diethyl ether.
`color{red}(2MH + B_2H_6 → 2 M^(+) [BH_4]^(–) \ \ \ \ (M = Li \ \ \ \ "or" \ \ \ \ Na))`
`color{green}(★)` Both `color{red}(LiBH_4)` and `color{red}(NaBH_4)` are used as reducing agents in organic synthesis. They are useful starting materials for preparing other metal borohydrides.