Tag Info

This tag is for questions concerning coordination compounds including but not limited to ligand properties, metal properties, orbital splitting, micro- and macroscopic properties of entire complexes etc. For complexes where carbon monoxide is the only ligand, use the [carbonyl-complexes] tag instead. For organic catalysts or proteins, the tag is applicable if the question is about the metal’s direct coordination sphere.

Coordination compounds are those that contain a dative or coordinative bond, usually formed by one atom donating a free lone pair to another, electron-deficient atom. In most cases, the electron deficient atom — a Lewis acid — is a (transition) metal atom while the donating atom is often a non-metal.

Every metal can more or less be considered to form coordination compounds in solution; for some, e.g. zinc, these are poorly defined and easily transition into each other. For others, e.g. aluminum or chromium, the ligand sphere is very stable and ligand exchanges require substantial activation energy. This is usually due to d-orbital splitting observed in complexes: While for a lone metal ion in a vacuum all five d-orbitals are degenerate if a set of ligands approaches said ion, the orbitals split up into sets depending on where the ligands approach from.

A large number of complexes is octahedral in structure. For those, the $\mathrm d_{xy}$, $\mathrm d_{xz}$ and $\mathrm d_{yz}$ (usually termed $\mathrm{t_{2g}}$ due to the way they transform in $O_\mathrm{h}$) remain unaltered at first approximation while the $\mathrm{e_g}$ orbitals $\mathrm d_{x^2 - y^2}$ and $\mathrm{d}_{z^2}$ — which point towards the ligands — are destabilised. This simple approximation is called crystal field approximation. For tetrahedral complexes, the $\mathrm{t_1}$ orbitals are destabilised and $\mathrm{e}$ orbitals slightly stabilised.

The simple crystal field model provides no quantitative approximation for the energy difference between the orbitals and thus for the light wavelength absorbed. If the energy difference is large enough, it may be thermodynamically favourable to keep all spins paired up in the lower energy levels generating a low-spin complex. If not, a high-spin complex is formed where the first five d-electrons are always filled in with parallel spins. Useful starting points for deciding whether a complex is high-spin or low spin is the spectrochemical series which orders ligands by the field split they typically cause. Halides are one extreme, usually termed weak field ligands since they almost always form high-spin complexes. Oxygen and nitrogen donors are generally medium field. The strong field end of the spectrum is formed by carbonyl and cyanide ligands which often exclusively form low-spin complexes such as $\ce{[Fe(CN)6]^4-}$. Note that the specialised tag exists for complexes containing entirely or predominantly $\ce{CO}$ as ligands.

Often, complexes are not perfect octahedra but distorted. This is often due to the Jahn-Teller effect: If the ligands on the $z$ axis are moved away slightly, all orbitals with a $z$ component are stabilised thus generating discreet energy levels rather than degenerate ones allowing for stabilisation of the ion’s state. Square planar complexes, which often occur for $\mathrm{d^8}$ systems, can be thought of as extreme Jahn-Teller cases where a pair of ligands has been completely removed from the system.

Many advanced organic reactions rely on catalysis by transition metal complexes, for example: