Introduction to Coordination Compounds

A coordination compound (or complex compound) is a compound in which a central metal atom or ion is bonded to a fixed number of surrounding ions or neutral molecules, called ligands, through coordinate (dative) covalent bonds — bonds in which the ligand alone supplies both electrons of the shared pair, while the metal supplies an empty orbital. The metal-plus-ligands unit is called the coordination entity (or coordination sphere), conventionally written inside square brackets, e.g. $[\text{Co}(\text{NH}_3)_6]^{3+}$ or $[\text{Fe}(\text{CN})_6]^{4-}$. Any ions needed to balance the overall charge of a charged coordination entity are written outside the brackets as counter ions — for example, in $[\text{Co}(\text{NH}_3)_6]\text{Cl}_3$, the complex cation $[\text{Co}(\text{NH}_3)_6]^{3+}$ is balanced by three $\text{Cl}^-$ counter ions.

This distinction between what is inside the coordination sphere (bonded directly to the metal, and therefore not free to react independently) and what is outside it (a genuinely free ion in solution) is the single most important structural idea in coordination chemistry, and it explains a great deal of otherwise puzzling chemical behaviour. A striking historical example: solid $\text{CoCl}_3.6\text{NH}_3$ reacts completely with excess $\text{AgNO}_3$ solution to precipitate all three chlorides as $\text{AgCl}$, while $\text{CoCl}_3.5\text{NH}_3$ precipitates only two of its three chlorides. Ordinary valence theory of the time could not explain why one chloride in the second compound behaved so differently from the other two — the explanation (developed fully in a later section on Werner's theory) is that in $[\text{Co}(\text{NH}_3)_5\text{Cl}]\text{Cl}_2$, one chloride has moved inside the coordination sphere, directly bonded to cobalt, and is no longer free to react as a simple ion.

Coordination compounds are chemically distinct from double salts. A double salt, such as carnallite ($\text{KCl}.\text{MgCl}_2.6\text{H}_2\text{O}$) or Mohr's salt ($\text{FeSO}_4.(\text{NH}_4)_2\text{SO}_4.6\text{H}_2\text{O}$), exists as a distinct crystalline solid but dissociates completely into its simple constituent ions when dissolved in water, and each of those ions gives its usual characteristic chemical test. A coordination compound, by contrast, retains its coordination entity largely intact in solution: dissolving $\text{K}_4[\text{Fe}(\text{CN})_6]$ gives $\text{K}^+$ ions and the intact complex ion $[\text{Fe}(\text{CN})_6]^{4-}$, which does not give the usual test for free $\text{Fe}^{2+}$ or free $\text{CN}^-$ ions — the cyanide is bound too strongly to the iron to behave as a separate species.

This chapter builds up the complete picture of coordination compounds: how ligands are classified (by denticity and by field strength), how coordination number determines shape, why complexes are coloured and magnetic in the specific ways they are, how they are systematically named, two complementary bonding theories (valence bond theory and crystal field theory) that explain their structure and magnetism, the several kinds of isomerism they display, and finally their genuine importance — in chemical analysis, in extracting metals like gold and silver, and in the biochemistry of life itself, from the oxygen-carrying haem group in blood to the cobalt at the heart of vitamin B12.