Chemistry Important trends and anomalous properties of boron and Some important compounds of boron

Topics covered

`=>` Important trends and anomalous properties of boron
`=>` Some important compounds of boron
`=>` Borax
`=>` Orthoboric acid
`=>` Diborane
`=>` Uses of boron and aluminium and their compounds


`color{green}(★)` Certain important trends can be observed in the chemical behaviour of group `13` elements.

`color{green}(★)` The tri-chlorides, bromides and iodides of all these elements being covalent in nature are hydrolysed in water.

`color{green}(★)` Species like tetrahedral `color{red}([M(OH)_4]^–)` and octahedral `color{red}([M(H_2O)_6]^(3+))`, except in boron, exist in aqueous medium.

`color{green}(★)` The monomeric trihalides, being electron deficient, are strong Lewis acids.

`color{green}(★)` Boron trifluoride easily reacts with Lewis bases such as `color{red}(NH_3)` to complete octet around boron.

`color{red}(F_3B + : NH_3 → F_3B ←NH_3)`

`color{green}(★)` It is due to the absence of `color{red}(d)` orbitals that the maximum covalence of `color{red}(B)` is `4`. Since the `color{red}(d)` orbitals are available with `color{red}(Al)` and other elements, the maximum covalence can be expected beyond `4`.

`color{green}(★)` Most of the other metal halides (e.g., `color{red}(AlCl_3)`) are dimerised through halogen bridging (e.g., `color{red}(Al_2Cl_6)`). The metal species completes its octet by accepting electrons from halogen in these halogen bridged molecules.

Q 3205001868

Boron is unable to form `BF_6^(3-)` ion. Explain.


Due to non-availability of d orbitals, boron is unable to expand its octet. Therefore, the maximum covalence of boron cannot
exceed 4.


Some useful compounds of boron are borax, orthoboric acid and diborane. We will briefly study their chemistry.


`color{green}(★)` It is the most important compound of boron. It is a white crystalline solid of formula `color{red}(Na_2B4O_7⋅10H_2O).`

`color{green}(★)` In fact it contains the tetranuclear units `color{red}([B_4O_5 (OH)_4]^(2-))` and correct formula; therefore, is `color{red}(Na_2 [ B_4O_5(OH)_4] . 8 H_2O)`. Borax dissolves in water to give an alkaline solution.

`color{red}(Na_2B_4O_7 + 7H_2O → 2NaOH + underset("Orthoboric acid")(4H_3BO_3))`

`color{green}(★)` On heating, borax first loses water molecules and swells up. On further heating it turns into a transparent liquid, which solidifies into glass like material known as borax bead.

`color{red}(Na_2B_4O_7. 10H_2O overset(Delta)→ underset("Sodium metaborate")(Na_2B_4O_7) overset(Delta)→ 2NaBO_2 + underset("Boric anhydride")(B_2O_3))`

`color{green}(★)` The metaborates of many transition metals have characteristic colours and, therefore, borax bead test can be used to identify them in the laboratory. For example, when borax is heated in a Bunsen burner flame with `color{red}(CoO)` on a loop of platinum wire, a blue coloured `color{red}(Co(BO_2)_2)` bead is formed.

Orthoboric acid

`color{green}(★)` Orthoboric acid, `color{red}(H_3BO_3)` is a white crystalline solid, with soapy touch. It is sparingly soluble in water but highly soluble in hot water. It can be prepared by acidifying an aqueous solution of borax.

`color{red}(Na_2B_4O_7 + 2HCl + 5H_2O → 2NaCl + 4B(OH)_3)`

`color{green}(★)` It is also formed by the hydrolysis (reaction with water or dilute acid) of most boron compounds (halides, hydrides, etc.).

`color{green}(★)` It has a layer structure in which planar `color{red}(BO_3)` units are joined by hydrogen bonds as shown in Fig. 11.1.

`color{green}(★)` Boric acid is a weak monobasic acid. It is not a protonic acid but acts as a Lewis acid by accepting electrons from a hydroxyl ion:

`color{red}(B(OH)_3 + 2HOH → [B(OH)_4]^(–) + H_3O^(+))`

`color{green}(★)` On heating, orthoboric acid above 370K forms metaboric acid, `color{red}(HBO_2)` which on further heating yields boric oxide, `color{red}(B_2O_3).`

`color{red}(H_3BO_3 overset(Delta)→ HBO_2 overset(Delta)→ B_2O_3)`

Q 3215001869

Why is boric acid considered as a weak acid?


Because it is not able to release `H^+` ions on its own. It receives `OH^–` ions from water molecule to complete its octet and in turn
releases `H^+` ions.

Diborane, `B_2H_6`

`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}(★)` Boron being extremely hard refractory solid of high melting point, low density and very low electrical conductivity, finds many applications.

`color{green}(★)` Boron fibres are used in making bullet-proof vest and light composite material for aircraft.

`color{green}(★)` The boron-10 (10B) isotope has high ability to absorb neutrons and, therefore, metal borides are used in nuclear industry as protective shields and control rods.

`color{green}(★)` The main industrial application of borax and boric acid is in the manufacture of heat resistant glasses (e.g., Pyrex), glass-wool and fibreglass.

`color{green}(★)` Borax is also used as a flux for soldering metals, for heat, scratch and stain resistant glazed coating to earthenwares and as constituent of medicinal soaps.

`color{green}(★)` An aqueous solution of orthoboric acid is generally used as a mild antiseptic.

`color{green}(★)` Aluminium is a bright silvery-white metal, with high tensile strength.

`color{green}(★)` It has a high electrical and thermal conductivity. On a weight-to-weight basis, the electrical conductivity of aluminium is twice that of copper.

`color{green}(★)` Aluminium is used extensively in industry and every day life.

`color{green}(★)` It forms alloys with `color{red}(Cu, Mn, Mg, Si)` and `color{red}(Zn)`.

`color{green}(★)` Aluminium and its alloys can be given shapes of pipe, tubes, rods, wires, plates or foils and, therefore, find uses in packing, utensil making, construction, aeroplane and transportation industry.

`color{green}(★)` The use of aluminium and its compounds for domestic purposes is now reduced considerably because of their toxic nature.