Unit 5: Coordination Compounds

Complex compounds — vital in biology, industry, and analytical chemistry

5.1 Werner's Theory

Alfred Werner (1893) proposed that metal ions exhibit two types of valences in coordination compounds:

        Primary valence: Ionisable, satisfied by negative ions (e.g., Cl⁻)

        Secondary valence: Non-ionisable, fixed for a metal, equal to coordination number

        The secondary valence determines the geometry: octahedral, tetrahedral, or square planar

Werner's theory explained why CoCl₃·6NH₃ (yellow) gives 3 mol AgCl, while CoCl₃·4NH₃ (green) gives only 1 mol AgCl — the chloride ions outside the coordination sphere are ionisable, those inside are not.

5.2 Key Terms

Coordination Entity

Central metal + ligands enclosed in square brackets. E.g., [Co(NH₃)₆]³⁺

Coordination Number

Number of donor atoms directly bonded to metal. Count only σ bonds.

Oxidation Number

Charge on metal if all ligands removed. E.g., Cu in [Cu(CN)₄]³⁻ → Cu(I)

Homoleptic vs Heteroleptic

Homoleptic: one kind of ligand. Heteroleptic: more than one kind.

5.3 IUPAC Nomenclature

        Key rules: Cation first → ligands alphabetically → central metal → oxidation state (Roman numeral)

        Anionic metal: suffix -ate (ferrate, cobaltate, platinate)

        Special ligand names: aqua (H₂O), ammine (NH₃), carbonyl (CO), nitrosyl (NO)

Examples

[Co(NH₃)₆]Cl₃ → hexaamminecobalt(III) chloride
K₃[Fe(CN)₆] → potassium hexacyanidoferrate(III)
[Cr(NH₃)₃(H₂O)₃]Cl₃ → triamminetriaquachromium(III) chloride

5.4 Isomerism

Geometrical Isomerism

Square planar [MA₂B₂]: cis (adjacent, 90°) and trans (opposite, 180°)
Octahedral [MA₂B₄]: cis and trans possible
Tetrahedral: No geometrical isomerism (all positions equivalent)

Structural Isomerism

Type	Description	Example
Linkage	Ambidentate ligand bonds through different atoms	[Co(NH₃)₅(NO₂)]²⁺ vs [Co(NH₃)₅(ONO)]²⁺
Coordination	Ligands swap between two metal centres	[Co(NH₃)₆][Cr(CN)₆] vs [Cr(NH₃)₆][Co(CN)₆]
Ionisation	Counter ion swaps with ligand	[Co(NH₃)₅Br]SO₄ vs [Co(NH₃)₅SO₄]Br
Solvate	Water inside vs outside coordination sphere	[Cr(H₂O)₆]Cl₃ vs [Cr(H₂O)₅Cl]Cl₂·H₂O

5.5 Valence Bond Theory

        VBT explains: Geometry, magnetic properties, and hybridization of coordination compounds

        Inner orbital complexes: d²sp³ (uses inner d orbitals, low spin, strong field ligands)

        Outer orbital complexes: sp³d² (uses outer d orbitals, high spin, weak field ligands)

Magnetic Moment

μ = √(n(n+2)) BM

Where n = number of unpaired electrons, BM = Bohr magneton. Each unpaired electron contributes ~1.73 BM.

5.6 Crystal Field Theory

        CFT basics: Ligands treated as point charges; d orbitals split into t₂g and eₓ sets

        Octahedral: Δₒ splitting (t₂g lower, eₓ higher)

        Tetrahedral: Δₜ = (4/9)Δₒ (inverted, smaller splitting)

        Colour: d-d transitions absorb visible light → complementary colour observed

Spectrochemical Series

I⁻ < Br⁻ < SCN⁻ < Cl⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < NH₃ < en < CN⁻ < CO

Weak field ligands (left) → small Δ → high spin complexes. Strong field ligands (right) → large Δ → low spin complexes.

Colour of Complexes

When a complex absorbs visible light, an electron is excited from t₂g to eₓ. The colour observed is complementary to the absorbed wavelength.

Complex	Absorbed (nm)	Colour Absorbed	Colour Observed
[Ti(H₂O)₆]³⁺	498	Blue-green	Violet
[Cu(H₂O)₄]²⁺	600	Red	Blue
[Co(NH₃)₆]³⁺	475	Blue	Yellow-orange

Anhydrous CuSO₄ is white (no ligand → no d-d transition). CuSO₄·5H₂O is blue because water ligands cause d-orbital splitting.

Summary

Coordination compounds contain a central metal ion bonded to ligands. Werner's theory introduced primary and secondary valences. IUPAC nomenclature provides systematic naming. Isomerism (geometrical and structural) explains different forms with the same formula. VBT explains geometry and magnetism through hybridization. CFT explains colour and magnetic properties through d-orbital splitting. The spectrochemical series ranks ligands by field strength. Coordination compounds are vital in biology (chlorophyll, haemoglobin), metallurgy, catalysis, and medicine.