Why atoms bond

Breaking one mole of H-H bonds takes 436 kJ, so two hydrogen atoms are far more stable together than apart.

Almost nothing around you is made of single, separate atoms. The air is mostly N₂ and O₂ molecules, water is H₂O, and a grain of salt is a block of sodium and chlorine packed together by the billion. Only the noble gases drift about as lone atoms. So the real question is not whether atoms join, but why they do, and why some combinations happen while others never do.

Lower energy is more stable

A ball on a hillside rolls down and stays in the valley. Atoms follow the same rule with energy. When two hydrogen atoms approach, the electron of each is attracted to both nuclei at once. Up to a point this attraction outweighs the repulsion between the two nuclei and between the two electrons, and the total energy falls. At a separation of 74 pm the energy reaches its lowest point. Push the nuclei closer and repulsion takes over; pull them apart and you must put energy back in. That valley is the bond.

The depth of the valley is the bond energy. For hydrogen it is 436 kJ for every mole of molecules, which is about 6 × 10²³ bonds. That energy has to be supplied to split the molecules, and it is released when the atoms join. Fuels work this way in reverse: when hydrogen burns, weaker bonds are broken and stronger ones in water are made, and the difference comes out as heat.

How can anyone measure a distance of 74 pm? Nobody uses a ruler. Molecules spin and vibrate, and they absorb light only at particular frequencies that depend on how far apart their atoms are and how stiff the bond is. Measure those frequencies with a spectrometer and the bond length and angle follow from the calculation. The NIST Computational Chemistry Comparison and Benchmark Database collects these measured geometries, and every bond length and angle in the 3D models in this course comes from it.

The full-shell pattern

Chemists have a quick way to predict which bonds will form. Noble gases, with full outer shells, barely react, so a full shell seems to be a comfortable arrangement. Many atoms bond in a way that gives each of them a full outer shell, usually 8 electrons, or 2 for hydrogen. This is called the octet rule. The American chemist G. N. Lewis described the shared electron pair in 1916, and his dot diagrams are still used to count electrons in bonds.

There are three ways to reach a full shell, and they give the three main kinds of bond. An atom can hand electrons to another atom, which makes ions held by electric attraction: an ionic bond. Two atoms can share pairs of electrons so that each counts the shared ones as its own: a covalent bond. Or a whole crowd of metal atoms can pool their outer electrons: a metallic bond. Which one happens depends mostly on electronegativity, the pull each atom has on electrons.

The octet rule is a guide, not a law. Boron in BF₃ ends up with only 6 outer electrons, and sulfur in SF₆ has 12. Transition metals follow their own counting. For the first twenty elements and the compounds of carbon, nitrogen and oxygen, though, it works remarkably often, and it is the right place to start.