Coordination Compounds
Easy Overview
Coordination compounds are everywhere — and I mean everywhere. The hemoglobin in your blood, the chlorophyll in plants, vitamin B₁₂, many industrial catalysts, and even the bright blue pigment Prussian blue are all coordination compounds. These are compounds where a central metal atom or ion is surrounded by a set of ligands (molecules or ions that donate electron pairs). They have fascinating structures, colors, and reactivities. Werner's coordination theory (1893) was revolutionary. Before Werner, chemists could not understand compounds like CoCl₃·6NH₃. Why does it behave like [Co(NH₃)₆]Cl₃ — three ionic chlorides? And why does CoCl₃·5NH₃ have only two ionic chlorides, with the third one bonded differently? Werner proposed that metals have primary valence (oxidation state — ionic bonds) and secondary valence (coordination number — dative bonds to ligands). The secondary valences point to fixed positions in space around the central atom — octahedral for Co(III), square planar for Ni(II). This was the birth of coordination chemistry. IUPAC nomenclature looks complicated but follows clear rules. Cation named first, then anion. Ligands in alphabetical order (ignoring prefixes like di-, tri-), followed by the metal name with its oxidation state in Roman numerals in parentheses. For anionic complexes, the metal name ends in -ate. Knowing nomenclature helps you decode the structure from the name and vice versa. The theories that explain bonding in coordination compounds — valence bond theory (VBT), crystal field theory (CFT), and ligand field theory (LFT) — build on each other to explain the geometry, magnetism, and colors of these compounds. VBT uses hybridization (sp³, dsp², d²sp³) to explain shape. CFT focuses on how ligands split d-orbital energies to explain color and magnetism. But CFT treats ligands as point charges — it is too simple. LFT adds more realistic covalent bonding. Together, these theories explain why some complexes are colored while others are colorless, why some are paramagnetic while others are diamagnetic, and why some are inert (slow to react) while others are labile (fast to react).
Werner's coordination theory
Alfred Werner proposed that metal atoms have two types of valences: primary valence (oxidation state — satisfied by anions, shown as charge on central atom) and secondary valence (coordination number — number of donor atoms bonded to the metal, shown by number of ligands). Primary valences are non-directional and satisfied by ionic bonds. Secondary valences are directional and point to fixed positions in space. In a series of cobalt-ammonia complexes: [Co(NH₃)₆]Cl₃ — all 6 NH₃ satisfy secondary valence, 3 Cl⁻ satisfy primary valence (all three Cl⁻ are ionic). [Co(NH₃)₅Cl]Cl₂ — 5 NH₃ + 1 Cl satisfy secondary valence, 2 Cl⁻ are ionic. [Co(NH₃)₄Cl₂]Cl — 4 NH₃ + 2 Cl in coordination sphere, 1 Cl⁻ ionic. [Co(NH₃)₃Cl₃] — all 3 Cl in coordination sphere, none ionic. Conductivity measurements confirm this: more ionic chlorides = higher conductivity. Werner assigned octahedral geometry to Co(III) complexes (6 secondary valences at 90° angles). He was awarded the Nobel Prize in 1913.
IUPAC nomenclature of coordination compounds
Rules: (1) Cation named before anion. (2) Within the coordination sphere: Ligands named first (alphabetical order, ignoring prefixes di-, tri-, tetra-, etc.), then the central metal. (3) Anionic ligands end in -o (chloride → chloro, cyanide → cyano, hydroxide → hydroxo, oxalate → oxalato). Neutral ligands keep their name (ammine for NH₃ — note double m, aqua for H₂O, carbonyl for CO). (4) If the complex is anionic, the metal name ends in -ate (ferrate, cuprate, cobaltate, nickelate). (5) Oxidation state of the metal is given in Roman numerals in parentheses. (6) Prefixes di-, tri-, tetra-, penta-, hexa- indicate number of each ligand. For complex ligands (like ethylenediamine = en), use bis-, tris-, tetrakis-. Examples: [Co(NH₃)₆]Cl₃ = hexaamminecobalt(III) chloride. K₄[Fe(CN)₆] = potassium hexacyanoferrate(II). [Pt(NH₃)₄][PtCl₄] = tetraammineplatinum(II) tetrachloroplatinate(II) — a coordination compound where both cation and anion are complexes! [Ni(CO)₄] = tetracarbonylnickel(0).
Valence Bond Theory (VBT)
Linus Pauling's VBT explains coordination using hybridization. The metal ion uses empty hybrid orbitals (s + p + d) to accept electron pairs from ligands. Major features: Coordination number 4 — two geometries possible: tetrahedral (sp³, e.g., [Ni(CO)₄], [Zn(NH₃)₄]²⁺) and square planar (dsp², e.g., [Ni(CN)₄]²⁻, [Pt(NH₃)₄]²⁺). Coordination number 6 — octahedral geometry (d²sp³, e.g., [Fe(CN)₆]⁴⁻, [Co(NH₃)₆]³⁺) or sp³d² (outer orbital, e.g., [FeF₆]³⁻, [CoF₆]³⁻). Inner orbital complexes (d²sp³) use (n-1)d orbitals and involve pairing of electrons (low spin, diamagnetic or fewer unpaired electrons). Outer orbital complexes (sp³d²) use nd orbitals (high spin, more unpaired electrons). [Co(NH₃)₆]³⁺: Co³⁺ (3d⁶), NH₃ is a strong field ligand → electrons pair in 3d, one 3d orbital empty → d²sp³ → inner orbital, diamagnetic. [CoF₆]³⁻: Co³⁺ (3d⁶), F⁻ is a weak field ligand → no pairing → 4 unpaired electrons → sp³d² → outer orbital, paramagnetic. VBT successfully predicts geometry and magnetic properties but fails to explain color and detailed spectra.
Crystal Field Theory (CFT) — splitting of d-orbitals
CFT treats ligands as point charges that repel d-electrons. In an octahedral field, the five d-orbitals split into two groups: t₂g (dxy, dyz, dzx) — lower energy, orbitals point between axes; eg (dx²-y², dz²) — higher energy, orbitals point directly at ligands along axes. The energy difference is Δ₀ (10 Dq). For octahedral complexes: t₂g set has 3 orbitals, each 0.4Δ₀ lower than the average energy; eg set has 2 orbitals, each 0.6Δ₀ higher. Δ₀ depends on the nature of the ligand (spectrochemical series) and the metal ion. Strong field ligands (CN⁻, CO) cause large splitting; weak field ligands (I⁻, Br⁻, F⁻, H₂O) cause small splitting. Spectrochemical series (increasing Δ): I⁻ < Br⁻ < S²⁻ < Cl⁻ < NO₃⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NH₃ < en < NO₂⁻ < CN⁻ < CO. Tetrahedral splitting (Δt) is smaller (~4/9 of Δ₀) and reversed in ordering: e set is lower, t₂ set is higher.
Crystal Field Theory — color and magnetism explained
For d⁴ to d⁷ configurations, two arrangements are possible depending on Δ relative to pairing energy (P). Weak field (small Δ, P > Δ): high spin — electrons occupy all five orbitals singly before pairing (Hund's rule). Strong field (large Δ, P < Δ): low spin — electrons pair up in t₂g before occupying eg. [Fe(CN)₆]⁴⁻: CN⁻ is strong field → large Δ → low spin: t₂g⁶ eg⁰ → 0 unpaired → diamagnetic. [Fe(H₂O)₆]²⁺: H₂O is weak field → small Δ → high spin: t₂g⁴ eg² → 4 unpaired → paramagnetic. The color of the complex is due to d-d transitions: an electron absorbs a photon of visible light and jumps from t₂g to eg. The energy difference (Δ) determines the wavelength absorbed and hence the complementary color observed. [Ti(H₂O)₆]³⁺ (d¹): t₂g¹ → Δ corresponds to green light → complex appears purple. [Cu(NH₃)₄]²⁺: NH₃ causes larger splitting → absorbs orange light → appears deep blue. CFT quantitatively explains the colors and magnetism of coordination compounds — something VBT could not do.
Colour in coordination compounds
Colour arises from electronic transitions between d-orbitals. When white light falls on a complex, it absorbs certain wavelengths to promote electrons from lower to higher d-orbitals. The transmitted (or reflected) light appears as the complementary colour. The absorbed wavelength is determined by Δ (splitting energy). Factors affecting Δ: (1) Nature of the ligand — stronger ligands (CN⁻, CO) produce larger Δ, absorb higher energy (shorter wavelength) light, shifting the complementary color. [Cu(H₂O)₆]²⁺ (pale blue — Δ ~13,600 cm⁻¹, absorbs red) vs [Cu(NH₃)₄]²⁺ (deep blue — Δ ~17,200 cm⁻¹, absorbs orange). (2) Nature of the metal ion — higher oxidation state → larger Δ (Fe²⁺ vs Fe³⁺). Heavier transition metals (3d vs 4d vs 5d) give larger Δ. (3) Geometry — tetrahedral complexes generally absorb at longer wavelengths (weaker splitting) and are less intensely colored. Selection rules: d-d transitions are Laporte-forbidden (parity forbidden) but become weakly allowed by vibronic coupling. This is why transition metal complexes are not intensely colored (ε ~ 10-100 L mol⁻¹ cm⁻¹). Charge transfer transitions (ligand → metal or metal → ligand) are allowed and produce very intense colors (ε ~ 10,000 L mol⁻¹ cm⁻¹) — e.g., the intense purple of KMnO₄ is due to O → Mn charge transfer, not d-d transitions.
Isomerism in coordination compounds
Coordination compounds exhibit structural isomerism (different atom connectivity) and stereoisomerism (same connectivity, different spatial arrangement). Structural isomerism: (1) Ionization isomerism — same formula, different ionizable anion outside the coordination sphere: [Co(NH₃)₅Br]SO₄ (maroon) vs [Co(NH₃)₅SO₄]Br (red). (2) Hydrate isomerism — water as ligand vs water of crystallization: [Cr(H₂O)₆]Cl₃ (violet), [Cr(H₂O)₅Cl]Cl₂·H₂O (green), [Cr(H₂O)₄Cl₂]Cl·2H₂O (dark green). (3) Linkage isomerism — ambidentate ligand can bind through different atoms: [Co(NH₃)₅NO₂]²⁺ (nitro, bonded through N, yellow) vs [Co(NH₃)₅ONO]²⁺ (nitrito, bonded through O, red). Stereoisomerism: (1) Geometrical (cis-trans) — in octahedral [Co(NH₃)₄Cl₂]⁺: cis isomer (violet, polar) vs trans isomer (green, nonpolar). In square planar [Pt(NH₃)₂Cl₂]: cisplatin (cis, anticancer drug) vs transplatin (trans, inactive). (2) Optical isomerism — non-superimposable mirror images. [Co(en)₃]³⁺ forms a pair of enantiomers (Δ and Λ forms). Optically active complexes rotate plane-polarized light. The biological activity of cisplatin and the inertness of transplatin is a classic example of how geometry determines therapeutic effect.
Bonding in metal carbonyls
Metal carbonyls (M-CO) are important in industrial catalysis and illustrate a unique bonding mode called synergic bonding. CO donates its lone pair on carbon to an empty orbital on the metal (σ-donation). Simultaneously, filled metal d-orbitals back-donate electron density into the empty π* (antibonding) orbital of CO (π-back donation). This back-donation strengthens the M-C bond and weakens the C-O bond (observed as a lower C-O stretching frequency in IR spectra — typically 1850-2100 cm⁻¹ vs 2143 cm⁻¹ for free CO). The more back-donation, the weaker the C-O bond. Terminal CO: one CO bonded to one metal (νCO ~2000-2100 cm⁻¹). Bridging CO: CO bonded to two or three metals (νCO ~1800-1900 cm⁻¹). Examples: Ni(CO)₄ (tetrahedral, Ni(0)), Fe(CO)₅ (trigonal bipyramidal), Cr(CO)₆ (octahedral). The 18-electron rule: stable organometallic complexes typically have 18 valence electrons (metal valence + electrons donated by ligands), mirroring the noble gas configuration. Ni(CO)₄: Ni(0) = 3d⁸4s² → 10 e⁻ + 4 × 2 e⁻ from CO = 18 e⁻. Fe(CO)₅: Fe(0) = 8 e⁻ + 5 × 2 = 18 e⁻. Cr(CO)₆: Cr(0) = 6 e⁻ + 6 × 2 = 18 e⁻. The 18-electron rule is analogous to the octet rule for main group elements.
Importance of coordination compounds
Biological systems: Hemoglobin — Fe²⁺ in a porphyrin ring coordinates O₂ reversibly. Myoglobin stores oxygen in muscles. Chlorophyll — Mg²⁺ in a porphyrin ring captures light energy for photosynthesis. Vitamin B₁₂ — Co³⁺ in a corrin ring is essential for red blood cell formation and nerve function. Many enzymes contain transition metals at their active sites — carbonic anhydrase (Zn), catalase (Fe), nitrogenase (Fe, Mo). Medicinal applications: Cisplatin [Pt(NH₃)₂Cl₂] is a potent anticancer drug — it binds to DNA, cross-links guanine bases, and prevents cell division. Since its discovery (1960s), platinum-based drugs have saved countless lives. In analytical chemistry: EDTA (ethylenediaminetetraacetic acid) is used in complexometric titrations to determine metal ion concentrations. Gravimetric analysis uses complex formation for selective precipitation. Industrial: Extraction of Ni and Co as carbonyls (Mond process). Purification of metals via complex formation. Tanning: Cr³⁺ complexes cross-link collagen fibers in leather production. Photography: AgBr complexes with thiosulfate (fixing). Coordination chemistry is the intersection of inorganic and biological chemistry with real-world applications.
Key Points
- •Werner's theory: primary valence (oxidation state) + secondary valence (coordination number)
- •IUPAC nomenclature: ligands alphabetical, anionic end in -o, anionic complex ends in -ate
- •VBT: hybridization (sp³, dsp², d²sp³, sp³d²) explains geometry and magnetism
- •Inner orbital (d²sp³, low spin) vs outer orbital (sp³d², high spin) complexes
- •CFT: d-orbital splitting in octahedral (t₂g + eg), tetrahedral (e + t₂), and square planar
- •Δ₀ depends on ligand (spectrochemical series) and metal ion
- •Weak field → high spin (more unpaired e⁻); strong field → low spin (fewer unpaired e⁻)
- •Colour: d-d transitions absorb specific λ → complementary colour observed
- •Isomerism: ionization, hydrate, linkage, geometrical (cis-trans), optical
- •Synergic bonding in metal carbonyls: σ-donation + π-back donation
- •18-electron rule: stable organometallics have 18 valence electrons
- •Importance: hemoglobin (Fe), chlorophyll (Mg), vitamin B₁₂ (Co), cisplatin (anticancer)
- •EDTA: hexadentate ligand used in complexometric titrations
Practice Questions
- Explain Werner's coordination theory. How did it explain the properties of CoCl₃·6NH₃ and CoCl₃·5NH₃?
- Write the IUPAC names of: (a) K₃[Fe(C₂O₄)₃] (b) [Co(NH₃)₅Cl]Cl₂ (c) [Ni(CO)₄].
- Explain crystal field theory. Draw the d-orbital splitting diagram for octahedral and tetrahedral complexes.
- What is the spectrochemical series? Explain high-spin and low-spin complexes with [Fe(CN)₆]⁴⁻ and [Fe(H₂O)₆]²⁺.
- What are the different types of isomerism in coordination compounds? Give one example of each.
- Explain synergic bonding in metal carbonyls. How does it affect the C-O bond?
- State and explain the 18-electron rule. Verify it for Ni(CO)₄ and Fe(CO)₅.
- Discuss the biological and medicinal importance of coordination compounds.