electron-deficient compounds

Most atoms in a Lewis structure end up surrounded by eight electrons. But some central atoms simply do not have enough valence electrons to go around, and they settle for fewer than an octet, or they share the few electrons they have very cleverly. These are electron-deficient compounds, and they are a speciality of boron and beryllium, near the left of the p-block and s-block.

The classic molecular example is boron trifluoride, BF3. Boron has only three valence electrons, so even after forming three B-F bonds it is surrounded by only six electrons — two short of an octet. That empty space makes BF3 a powerful Lewis acid: it is hungry to accept a lone pair from something else, forming an adduct such as F3B-NH3. An even more striking case is the boron hydrides (boranes) like diborane B2H6, where there are too few electrons to draw an ordinary structure at all; the molecule solves the problem with three-centre two-electron bonds, in which a single pair of electrons holds together a B-H-B bridge spanning three atoms.

Electron-deficient compounds matter because their hunger for electrons makes them reactive and useful: boron trihalides are go-to Lewis-acid catalysts, and the strange bonding of boranes and carboranes spawned a whole sub-field of cluster chemistry summarised by Wade's rules. They are also a clean reminder that the octet rule is only a guideline — sometimes nature genuinely makes do with less.