Outer Orbital Complex
Definition and meaning of Outer Orbital Complex in chemistry.
An outer orbital complex is a coordination compound in which the d-orbitals used in bonding originate from the outermost d-subshell of the central metal ion. These complexes are typically high-spin and paramagnetic.
In more detail
In octahedral complexes, outer orbital bonding involves sp3d2 hybridization. These complexes form with weak-field ligands such as fluoride, chloride, or water, which produce relatively small ligand field splitting. The modest splitting energy allows electrons to remain unpaired, following Hund's rule rather than pairing in lower-energy orbitals. This contrasts sharply with inner orbital complexes, which employ d2sp3 hybridization with strong-field ligands and typically form low-spin configurations that are diamagnetic in common d6 cases (e.g., Co(III) or Fe(II) hexaammine complexes) but can retain unpaired electrons for other d-electron counts (e.g., d5 or d7). Although this terminology is older, it effectively classifies coordination compounds by their magnetic properties and electronic structure.
Key facts
| Hybridization | sp3d2 |
|---|---|
| Spin State | High-spin, paramagnetic |
| Ligand Field | Weak-field ligands (F-, Cl-, H2O) |
| Field | Inorganic Chemistry |
[FeF6]3- is a classic outer orbital complex. The iron(III) ion (d5 configuration) retains five unpaired electrons when coordinated by weak-field fluoride ligands, using sp3d2 hybridization and exhibiting paramagnetism.
Frequently asked questions
How do outer orbital and inner orbital complexes differ?
Outer orbital complexes use outermost d-orbitals with sp3d2 hybridization and form high-spin, paramagnetic structures with weak-field ligands. Inner orbital complexes use d2sp3 hybridization with strong-field ligands to form low-spin structures, which are diamagnetic in common d6 cases but can still have unpaired electrons for other d-electron counts.
Why are outer orbital complexes typically paramagnetic?
Weak-field ligands produce small ligand field splitting, so the energy cost of pairing electrons exceeds the orbital energy difference. Electrons therefore occupy separate orbitals with parallel spins, creating net magnetic moment.