boron trihalide Lewis acids (BF3, BCl3)

Imagine a molecule walking around with an empty seat at its table, perpetually inviting a guest to sit down. The boron trihalides, BF3, BCl3, BBr3, are exactly that: small, flat, electron-hungry molecules that crave a lone pair of electrons. They are among the cleanest examples of a Lewis acid, an electron-pair acceptor, that you will meet.

In a boron trihalide BX3 the boron is bonded to three halogens in a trigonal-planar shape with 120-degree angles, leaving boron with only six valence electrons and a completely empty p orbital sticking up perpendicular to the plane. That empty orbital is the open seat: any species with a lone pair (ammonia, ether, fluoride ion) can donate into it, forming an adduct in which boron becomes four-coordinate and tetrahedral, for example F3B-NH3 or [BF4]-. Curiously, the Lewis acidity does not follow electronegativity simply: BF3 is a weaker Lewis acid than BCl3 and BBr3, because in BF3 the small fluorine lone pairs donate back into boron's empty p orbital (pi back-bonding), partly satisfying it; the heavier halogens overlap poorly with boron, leave the p orbital emptier, and so make a stronger acid.

These molecules matter as practical catalysts and reagents. BF3 (often used as its diethyl ether complex) catalyzes many organic reactions by activating carbonyls and generating cations; BCl3 and BBr3 cleave ethers and are used in semiconductor processing. They are textbook anchors for Lewis acid-base theory, the empty-orbital picture, and the surprising role of pi back-bonding in tuning acidity. They are also corrosive and fume in moist air (hydrolyzing to boric acid and the hydrogen halide), so they are handled with care.