Coordination Compounds — Notes

Double salts (Mohr’s salt, carnallite) give all their simple ions in water; coordination compounds keep the complex ion intact ($\ce{K4[Fe(CN)6]}$ gives no free $\ce{Fe^2+}$ or $\ce{CN-}$).

Werner’s theory: a metal shows primary valency (ionisable; equals the oxidation state; satisfied by anions) and secondary valency (non-ionisable; equals the coordination number; satisfied by ligands; directed in space). Silver nitrate precipitates only chloride outside the coordination sphere: $\ce{CoCl3.6NH3}$ → 3 AgCl, $\ce{[Co(NH3)6]Cl3}$; $\ce{CoCl3.5NH3}$ → 2, $\ce{[CoCl(NH3)5]Cl2}$; $\ce{CoCl3.4NH3}$ → 1, $\ce{[CoCl2(NH3)4]Cl}$.

Terms: coordination entity, central atom/ion, ligand (electron-pair donor), coordination number (number of donor atoms bonded), coordination sphere, counter ions, coordination polyhedron, oxidation number, homoleptic and heteroleptic complexes.

Ligands: unidentate ($\ce{Cl-}$, $\ce{H2O}$, $\ce{NH3}$), didentate (en, oxalate), polydentate (EDTA⁴⁻, hexadentate). A chelate forms when a di- or polydentate ligand closes a ring with the metal; chelates are extra stable. Ambidentate ligands bind through either of two atoms: $\ce{NO2-}$/$\ce{ONO-}$, $\ce{SCN-}$/$\ce{NCS-}$.

1. Name the cation first, then the anion.
2. Inside the complex, name ligands in alphabetical order (ignoring di-, tri-, bis-, tris-), then the metal.
3. Prefixes: di-, tri-, tetra-, …; for ligands whose names contain a number (ethane-1,2-diamine), use bis-, tris-, tetrakis- with brackets.
4. Anionic ligands end in -o/-ido (chlorido, cyanido, hydroxido, oxalato); neutral ligands keep their names except aqua, ammine, carbonyl, nitrosyl.
5. Give the metal’s oxidation state in Roman numerals.
6. In an anionic complex the metal ends in -ate (ferrate, cuprate, argentate, aurate, plumbate, stannate, cobaltate, nickelate).
7. Ambidentate ligands: nitrito-N / nitrito-O, thiocyanato-S / thiocyanato-N.

Examples: $\ce{[CoCl2(NH3)4]Cl}$ tetraamminedichloridocobalt(III) chloride; $\ce{K3[Fe(CN)6]}$ potassium hexacyanidoferrate(III); $\ce{[Ni(CO)4]}$ tetracarbonylnickel(0); $\ce{[Co(en)3]Cl3}$ tris(ethane-1,2-diamine)cobalt(III) chloride; $\ce{K3[Al(C2O4)3]}$ potassium trioxalatoaluminate(III).

Structural isomerism:

• Ionisation: ligand and counter ion interchange — $\ce{[Co(NH3)5(SO4)]Br}$ / $\ce{[CoBr(NH3)5]SO4}$.
• Linkage: ambidentate ligand bonds through a different atom — nitrito-N / nitrito-O.
• Coordination: ligands interchange between cationic and anionic entities — $\ce{[Co(NH3)6][Cr(CN)6]}$ / $\ce{[Cr(NH3)6][Co(CN)6]}$.
• Solvate (hydrate): water inside or outside the sphere — $\ce{[Cr(H2O)6]Cl3}$ (violet) / $\ce{[CrCl(H2O)5]Cl2.H2O}$ (grey-green).

Stereoisomerism:

• Geometrical: square planar $\ce{[MX2L2]}$ and octahedral $\ce{[MX2L4]}$ give cis and trans; octahedral $\ce{[MA3B3]}$ gives fac and mer. Tetrahedral complexes show none.
• Optical: non-superimposable mirror images (d and l), e.g. $\ce{[Co(en)3]^3+}$ and cis-$\ce{[CoCl2(en)2]+}$; trans-$\ce{[CoCl2(en)2]+}$ is optically inactive.

VBT: the metal ion uses hybrid orbitals to accept ligand electron pairs. CN 4: $sp^3$ tetrahedral ($\ce{[NiCl4]^2-}$, paramagnetic) or $dsp^2$ square planar ($\ce{[Ni(CN)4]^2-}$, diamagnetic). CN 6: $d^2sp^3$ inner-orbital/low spin ($\ce{[Co(NH3)6]^3+}$, diamagnetic) or $sp^3d^2$ outer-orbital/high spin ($\ce{[CoF6]^3-}$, 4 unpaired). Limitations: no explanation of colour; cannot predict tetrahedral vs square planar; no distinction of weak and strong ligands.

CFT: ligands are point charges. Octahedral: $e_g$ ($d_{x^2-y^2}$, $d_{z^2}$) at $+0.6\Delta_o$, $t_{2g}$ at $-0.4\Delta_o$. Tetrahedral: e lower, $t_2$ higher, $\Delta_t \approx \tfrac{4}{9}\Delta_o$ — always high spin.

Spectrochemical series: $\ce{I-} < \ce{Br-} < \ce{SCN-} < \ce{Cl-} < \ce{S^2-} < \ce{F-} < \ce{OH-} < \ce{C2O4^2-} < \ce{H2O} < \ce{NCS-} < \text{edta}^{4-} < \ce{NH3} < \text{en} < \ce{CN-} < \ce{CO}$. If $\Delta_o < P$: high spin ($d^4$: $t_{2g}^3e_g^1$); if $\Delta_o > P$: low spin ($t_{2g}^4e_g^0$).

Colour: d–d transitions absorb visible light; the complementary colour is seen ($\ce{[Ti(H2O)6]^3+}$ absorbs about 498 nm and looks violet; $\ce{[Ni(H2O)6]^2+}$ green → $\ce{[Ni(en)3]^2+}$ violet). Anhydrous $\ce{CuSO4}$ is white. Limitation of CFT: anionic ligands should split most, yet halides are weak; covalent bonding is ignored.

Metal carbonyls: $\ce{[Ni(CO)4]}$ tetrahedral, $\ce{[Fe(CO)5]}$ trigonal bipyramidal, $\ce{[Cr(CO)6]}$ octahedral; $\ce{[Mn2(CO)10]}$ and $\ce{[Co2(CO)8]}$ have metal–metal bonds. Synergic bonding: σ donation from the carbon lone pair into an empty metal orbital, plus π back-donation from a filled metal d orbital into the empty π* orbital of CO; the M–C bond is strengthened and the C–O bond weakened.

Applications: estimation of hardness of water with EDTA ($\ce{Ca^2+}$, $\ce{Mg^2+}$); extraction of silver and gold as cyanide complexes; purification of nickel by the Mond process via $\ce{[Ni(CO)4]}$; Wilkinson’s catalyst $\ce{[RhCl(PPh3)3]}$ for hydrogenation; photography (AgBr dissolved by hypo as $\ce{[Ag(S2O3)2]^3-}$); medicine (cisplatin against cancer; EDTA to remove lead); biology (chlorophyll Mg, haemoglobin Fe, vitamin B12 Co).