Transition metals and coordination compounds

Main-group metals like sodium lose their few outermost electrons and stop, landing on one stable oxidation state. Transition metals are built differently: their outer s electrons and an inner, incompletely filled d subshell sit at similar energies, so a transition metal can lose different numbers of electrons for only a modest energy difference. Iron loses two electrons to become Fe²⁺ or three to become Fe³⁺, both readily accessible — variable oxidation state is the group's defining feature and follows directly from that near-degeneracy.

Coordination chemistry follows from the same d subshell. A transition metal ion has empty or partially filled d orbitals that can accept electron pairs donated by surrounding molecules or ions, called ligands — water, ammonia, chloride, cyanide. The resulting coordination complex, metal plus attached ligands, is held together by bonds that are electronically just ordinary covalent bonds, distinguished only by both electrons in the shared pair originating from the ligand rather than one from each partner.

The colours associated with transition metal compounds come from exactly this arrangement. Surrounding ligands split the d orbitals' energies apart by an amount that depends on which ligands are present, and electrons absorb visible light jumping between the split levels — the wavelength absorbed, and so the colour seen, shifts with the ligand. This is why the same metal ion, copper say, gives a pale blue solution with water as ligand and a much deeper blue with ammonia: nothing about the copper has changed, only what surrounds it and, consequently, the size of the gap the light must bridge.