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Boron trifluoride is the inorganic compound with the formula BF 3. This pungent, colourless, and toxic gas forms white fumes in moist air. It is a useful Lewis acid and a versatile building block for other boron compounds.
Lewis dot diagram structures show three formal alternatives for describing bonding in boron monofluoride. BF is unusual in that the dipole moment is inverted with fluorine having a positive charge even though it is the more electronegative element. This is explained by the 2sp orbitals of boron being reoriented and having a higher electron density.
[1] [2] [3] Introduced by Gilbert N. Lewis in his 1916 article The Atom and the Molecule, a Lewis structure can be drawn for any covalently bonded molecule, as well as coordination compounds. [4] Lewis structures extend the concept of the electron dot diagram by adding lines between atoms to represent shared pairs in a chemical bond.
Some of the most studied examples of such Lewis acids are the boron trihalides and organoboranes: [9] BF 3 + F − → BF − 4. In this adduct, all four fluoride centres (or more accurately, ligands) are equivalent. BF 3 + OMe 2 → BF 3 OMe 2. Both BF 4 − and BF 3 OMe 2 are Lewis base adducts of boron trifluoride.
The trifluoride is produced by treating borate salts with hydrogen fluoride, while the trichloride is produced by carbothermic reduction of boron oxides in the presence of chlorine gas: [49] [51] B 2 O 3 + 3 C + 6 Cl 2 → 2 BCl 3 + 3 CO Boron (III) trifluoride structure, showing "empty" boron p orbital in pi-type coordinate covalent bonds
The trihalides adopt a planar trigonal structure. These compounds are Lewis acids in that they readily form adducts with electron-pair donors, which are called Lewis bases. For example, fluoride (F −) and boron trifluoride (BF 3) combined to give the tetrafluoroborate anion, BF 4 −. Boron trifluoride is used in the petrochemical industry as ...
Bonding diagram of diborane (B 2 H 6) showing with curved lines a pair of three-center two-electron bonds, each of which consists of a pair of electrons bonding three atoms; two boron atoms and a hydrogen atom in the middle. The structure of diborane has D 2h symmetry. Four hydrides are terminal, while two bridge between the boron centers.
The boron atom in BH 3 has 6 valence electrons. Consequently, it is a strong Lewis acid and reacts with any Lewis base ('L' in equation below) to form an adduct: [7] BH 3 + L → L—BH 3. in which the base donates its lone pair, forming a dative covalent bond. Such compounds are thermodynamically stable, but may be easily oxidised in air.