1. Werner's Theory of Coordination Compounds

Alfred Werner proposed the concept of primary and secondary valencies for metal ions in coordination compounds.

• Primary Valency: Corresponds to the oxidation state of the central metal ion. It is non-directional, ionisable (can be precipitated), and satisfied by negative ions.
• Secondary Valency: Corresponds to the coordination number of the central metal ion. It is directional, non-ionisable, and satisfied by negative ions or neutral molecules (ligands). Every metal has a fixed number of secondary valencies.

1.1 Double Salts vs Complex Salts

• Double Salts: Dissociate completely into simple ions when dissolved in water. Example: Mohr's salt [FeSO₄·(NH₄)₂SO₄·6H₂O], Potash alum [KAl(SO₄)₂·12H₂O].
• Complex Salts: Contain a complex ion that does not dissociate into simple ions in water. Example: K₄[Fe(CN)₆] yields 4K⁺ and [Fe(CN)₆]⁴⁻ ions, but NOT Fe²⁺ or CN⁻ ions.

2. Definitions of Some Important Terms

• Central Atom/Ion: The atom/ion to which a fixed number of ions/groups are bound in a definite geometrical arrangement around it. It acts as a Lewis acid (electron pair acceptor).
• Ligands: The ions or molecules bound to the central atom/ion. They act as Lewis bases (electron pair donors).
  • Unidentate: Donates one electron pair (e.g., Cl⁻, H₂O, NH₃).
  • Didentate / Bidentate: Donates two electron pairs (e.g., ethane-1,2-diamine 'en', oxalate ion C₂O₄²⁻ 'ox').
  • Polydentate: Donates several electron pairs (e.g., EDTA⁴⁻ is hexadentate).
  • Chelate Ligand: A di- or polydentate ligand that uses its two or more donor atoms to bind a single metal ion, forming a ring structure. Chelate complexes are more stable.
  • Ambidentate Ligand: A ligand which can ligate through two different atoms. (e.g., NO₂⁻ can bind through N or O; SCN⁻ can bind through S or N).
• Coordination Number (C.N.): The number of ligand donor atoms to which the metal is directly bonded. (Note: A bidentate ligand counts as 2 towards C.N.).
• Coordination Polyhedron: The spatial arrangement of the ligand atoms around the central atom (e.g., octahedral, tetrahedral, square planar).

3. IUPAC Nomenclature of Coordination Compounds

Rules for naming mononuclear coordination compounds:

1. The cation is named first in both positively and negatively charged coordination entities.
2. The ligands are named in an alphabetical order before the name of the central atom/ion.
3. Names of anionic ligands end in -o (e.g., chlorido, cyanido). Neutral ligands use their common names, but there are exceptions: H₂O is aqua, NH₃ is ammine, CO is carbonyl, NO is nitrosyl.
4. Prefixes di-, tri-, tetra-, etc., are used to indicate the number of individual ligands. If the ligand name itself includes a numerical prefix (e.g., ethylenediamine), the terms bis-, tris-, tetrakis- are used, and the ligand name is placed in parentheses.
5. If the complex ion is an anion, the name of the metal ends with the suffix -ate (e.g., ferrate for iron, argentate for silver). If the complex ion is a cation or neutral, the metal name is the same as the element.
6. The oxidation state of the metal in cation, anion or neutral coordination entity is indicated by Roman numeral in parentheses.

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

4. Isomerism in Coordination Compounds

Isomers have the same molecular formula but different arrangements of atoms. They differ in physical and/or chemical properties.

4.1 Structural Isomerism

• Linkage Isomerism: Arises in a coordination compound containing an ambidentate ligand. E.g., [Co(NH₃)₅(NO₂)]Cl₂ (with NO₂ bonded through N) and [Co(NH₃)₅(ONO)]Cl₂ (bonded through O).
• Coordination Isomerism: Arises from the interchange of ligands between cationic and anionic entities of different metal ions present in a complex. E.g., [Co(NH₃)₆][Cr(CN)₆] and [Cr(NH₃)₆][Co(CN)₆].
• Ionisation Isomerism: E.g., [Co(NH₃)₅SO₄]Br vs [Co(NH₃)₅Br]SO₄. They give different ions in solution.
• Solvate (Hydrate) Isomerism: E.g., [Cr(H₂O)₆]Cl₃ (violet) vs [Cr(H₂O)₅Cl]Cl₂·H₂O (grey-green).

4.2 Stereoisomerism

• Geometrical Isomerism: Arises in heteroleptic complexes due to different possible geometric arrangements of the ligands.
  • Square Planar (ML₂X₂): Can exist as cis (similar ligands adjacent to each other) or trans (opposite to each other). Examples: [Pt(NH₃)₂Cl₂] (cis-platin is used in cancer therapy).
  • Octahedral (ML₄X₂): Can exist as cis and trans.
  • Octahedral (ML₃X₃): Forms fac (facial) and mer (meridional) isomers.
• Optical Isomerism: Chiral molecules that are non-superimposable mirror images of each other. They are called enantiomers (dextro and laevo forms). Common in octahedral complexes involving didentate ligands like [Co(en)₃]³⁺ or cis-[PtCl₂(en)₂]²⁺. Trans isomers are usually optically inactive due to a plane of symmetry.

5. Valence Bond Theory (VBT)

Predicts the shape and magnetic properties based on hybridization.

• Coordination Number 4:
  • sp³ Hybridization: Tetrahedral shape. Generally with weak field ligands (e.g., [NiCl₄]²⁻).
  • dsp² Hybridization: Square planar shape. Generally with strong field ligands that pair up inner d-electrons (e.g., [Ni(CN)₄]²⁻).
• Coordination Number 6:
  • d²sp³ Hybridization: Inner orbital complex. Shape: Octahedral. Occurs when inner (n-1)d orbitals are used (usually with strong field ligands). E.g., [Co(NH₃)₆]³⁺.
  • sp³d² Hybridization: Outer orbital complex. Shape: Octahedral. Occurs when outer nd orbitals are used (usually with weak field ligands). E.g., [CoF₆]³⁻.

6. Crystal Field Theory (CFT)

Considers the metal-ligand bond to be completely ionic (electrostatic interaction). When ligands approach the central metal ion, the five degenerate d-orbitals split into two sets of different energies (Crystal Field Splitting).

6.1 Crystal Field Splitting in Octahedral Complexes

The six ligands approach along the x, y, and z axes.

• The orbitals lying strictly on the axes (d_x²-y², d_z²) experience more repulsion and their energy is raised. This set is called e_g.
• The orbitals lying between the axes (d_xy, d_yz, d_zx) experience less repulsion and their energy is lowered. This set is called t_2g.
• The energy difference between t_2g and e_g is called crystal field splitting energy, denoted by Δ_o.

6.2 Crystal Field Splitting in Tetrahedral Complexes

The four ligands approach between the axes.

• The d_xy, d_yz, d_zx (the t₂ set) orbitals are raised in energy.
• The d_x²-y², d_z² (the e set) orbitals are lowered in energy.
• Splitting energy is smaller. Δ_t = (4/9) Δ_o. Because Δ_t is small, pairing does not usually occur (mostly high spin complexes).

6.3 Spectrochemical Series

Arrangement of ligands in order of increasing field strength:
I⁻ < Br⁻ < SCN⁻ < Cl⁻ < S²⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < edta⁴⁻ < NH₃ < en < CN⁻ < CO

• Weak field ligands (e.g., Halogens): Cause small splitting (Δ_o < P, pairing energy). Electrons do not pair up in t_2g but enter e_g. Form High spin complexes.
• Strong field ligands (e.g., CN⁻, CO): Cause large splitting (Δ_o > P). Electrons pair up in t_2g before filling e_g. Form Low spin complexes.

6.4 Colour in Coordination Compounds

Colour is due to the excitation of an electron from lower energy d-orbital (e.g., t_2g) to a higher energy d-orbital (e.g., e_g). This is called d-d transition.

If ligands are removed, splitting disappears, and the substance becomes colourless (e.g., when anhydrous CuSO₄ loses water, it becomes white).

7. Bonding in Metal Carbonyls

The metal-carbon bond in metal carbonyls possess both s and p character. The ligand CO is a pi-acid ligand.

This electron transfer from metal to ligand is called back bonding. This creates a synergic effect which strengthens the bond between CO and the metal.

🎓 NEET Previous Year Questions

📝 IUPAC Nomenclature (Q1 – Q6)

💡 VBT & CFT Reasoning (Q7 – Q14)

📈 Isomerism & Applications (Q15 – Q22)