lone-pair repulsion

Methane and water both have four electron pairs around their central atom, so why is methane's bond angle a perfect 109.5 degrees while water's is squeezed down to about 104.5 degrees? The answer is lone-pair repulsion: a lone pair pushes harder on its neighbours than a bonding pair does, bending the rest of the molecule out of the textbook angles.

The reason is a matter of how the electron pairs are held. A bonding pair is shared between two nuclei, so it is pulled tight into the bond, like a rope held at both ends. A lone pair is held by only one nucleus, so it billows out, fatter and closer to the central atom, occupying more angular space. The result is a clear pecking order of repulsion strength: lone pair to lone pair is strongest, lone pair to bonding pair is intermediate, and bonding pair to bonding pair is weakest. Lone pairs therefore claim the roomier positions and compress the bond angles of everything around them.

This single idea fine-tunes VSEPR predictions across inorganic chemistry. It explains the angle order methane > ammonia > water (zero, one, then two lone pairs on the central atom, each shrinking the angle further); why in a trigonal bipyramid lone pairs always take the roomier equatorial positions (as in SF4, ClF3, XeF2); and why an octahedron with two lone pairs ends up square planar with the lone pairs trans to each other, as in XeF4. Without lone-pair repulsion, VSEPR would predict the wrong shape for a great many real molecules.