Q.Describe the shapes of BF3 and BH4–. Assign the hybridisation of boron in these species.

Diborane, $B_2H_6$, is the simplest boron hydride: a colourless, highly toxic gas that ignites spontaneously in air, burning with a very large release of energy ($\Delta_cH^\circ = -1976$ kJ/mol), and it hydrolyses readily to boric acid. It can be made in the lab from $BF_3$ with $LiAlH_4$ or by oxidising $NaBH_4$ with iodine, and industrially from $BF_3$ and $NaH$ at 450 K. Structurally, the two borons and four terminal hydrogens lie in one plane, with two bridging hydrogens above and below; the four terminal B–H bonds are ordinary two-centre-two-electron bonds, but the two bridging B–H–B bonds are unusual three-centre-two-electron ("banana") bonds. Each boron is $sp^3$ hybridised, and of its four hybrid orbitals, one is left electron-less and used for bridge bonding — this is the structural root of diborane's electron deficiency. Diborane cleaves with Lewis bases (L) to give $BH_3 \cdot L$ adducts (with $NMe_3$, CO), and reacts differently with ammonia — first forming the ionic $[BH_2(NH_3)_2]^+[BH_4]^-$, which on heating converts to borazine ($B_3N_3H_6$, "inorganic benzene"). Diborane also underlies the chemistry of the tetrahedral borohydride ion $[BH_4]^-$: $LiBH_4$ and $NaBH_4$, made by reacting the metal hydride with $B_2H_6$, are important reducing agents in organic synthesis. Separately, boron's other hydride-derived species — $BF_3$ and $BH_4^-$ — illustrate contrasti …