Chemical bonding

Atoms bond because the bonded arrangement is lower in energy than the separated one. That single fact settles the most persistent misconception in chemistry: breaking a bond always costs energy, and forming one always releases it. Reactions give out heat when the bonds formed in the products are stronger than the bonds broken in the reactants — never because breaking bonds released something.

The textbook categories are ionic and covalent: electrons transferred outright versus electrons shared. But these are the two ends of one continuum, not two separate species. The controlling quantity is electronegativity difference. A large difference pulls the electron across completely and gives ions; no difference gives equal sharing; anything between gives a polar covalent bond, shared but lopsided. Water is the important middle case — the O–H bonds are covalent yet strongly polar, which is what makes water an excellent solvent for salts and the reason ice floats.

That position on the continuum is what actually predicts behaviour. Ionic solids form rigid lattices, so they melt high, dissolve in water and conduct once molten or dissolved. Covalent molecular substances are held to each other only by weak intermolecular forces, so they melt low and rarely conduct. The octet rule is a useful pattern behind much of this, but it is a rule of thumb with real exceptions — boron is content with six electrons, phosphorus and sulfur routinely exceed eight — and treating it as a law causes more errors than it prevents.