Revision: Coordination Compounds Chemistry HSC Science (General) 12th Standard Board Exam Maharashtra State Board

Anomers are stereoisomers of sugars that differ only in the configuration of the hydroxyl group at the C₁ position. They are not mirror images of each other.

Alfred Werner (1893) proposed the first systematic theory to explain the structure and bonding in coordination compounds. His key postulates are:

Main Postulates:

1. In coordination compounds, metals show two types of valencies — Primary (ionisable) and Secondary (non-ionisable).
2. Primary valences are normally ionisable. They are satisfied by negative ions (counter ions/anions). They correspond to the metal's oxidation state.
3. Secondary valences are non-ionisable. They are satisfied by neutral molecules or negative ions (ligands). The secondary valency is equal to the coordination number and is constant for a metal.
4. Ion groups bound by secondary valencies to the metal have a characteristic spatial arrangement (geometry). This geometry is decided by the secondary valences, not the primary valences.

Werner's Formula Examples:

• CrCl₃·6H₂O: In [Cr(H₂O)₆]Cl₃, all three Cl⁻ are outside the coordination sphere and hence ionisable.
• CrCl₃·5H₂O: [Cr(H₂O)₅Cl]Cl₂ — two Cl⁻ are ionisable, one is inside the coordination sphere.
• CrCl₃·4H₂O: [Cr(H₂O)₄Cl₂]Cl — one Cl⁻ ionisable.
• CrCl₃·3H₂O: [Cr(H₂O)₃Cl₃] — no ionisable Cl⁻, no precipitate with AgNO₃.

• Proposed by Heitler and London (1927), further developed by Pauling and Slater.
• A covalent bond is formed when half-filled valence atomic orbitals of similar energies overlap, each containing one unpaired electron.
• Greater the overlap → stronger the bond.

Types of Orbital Overlap:

Hybridisation & Shapes:

Limitations of VBT:

• Involves a number of assumptions.
• Does not give a quantitative interpretation of magnetic data.
• Does not explain the colour exhibited by coordination compounds.
• Does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds.
• Does not make exact predictions regarding the tetrahedral and square planar structures of 4-coordinate complexes.
• Does not distinguish between weak and strong ligands.

Ligands:

Ligands are the donor atoms, molecules, or anions that donate a pair of electrons to the metal atom or ion and form a coordinate bond. The number of coordinating or ligating groups present in a ligand is called the denticity of that ligand.

Coordination Number:

The coordination number (CN) of a metal ion in a complex is the total number of unidentate ligands (plus double the number of didentate ligands if any) attached to the central metal ion through coordinate bonds.

Homoleptic vs Heteroleptic Complexes:

1. Homoleptic complexes: Metal is bound to only one kind of donor group. e.g., [Co(NH₃)₆]³⁺
2. Heteroleptic complexes: Metal is bound to more than one kind of donor group. e.g., [Co(NH₃)₄Cl₂]⁺

• In coordination compounds, d-orbitals split into t₂g (lower) and e_g (higher) energy levels due to the ligand field.
• The energy difference between them is called the crystal field splitting energy (Δ₀).
• This Δ₀ lies in the visible region, so these compounds absorb visible light.
• When light is absorbed, an electron jumps from t₂g → e_g, called a d–d transition.
• The observed colour is complementary to the colour of light absorbed.
• The energy relation is: E = hν = Δ₀.
• Metal ions with d¹–d⁹ configuration are coloured, while d⁰ and d¹⁰ are colourless.
• Some compounds (e.g., KMnO₄) show colour due to charge transfer (LMCT), not d–d transition.
• Ligand strength affects colour: strong field ligands ↑ Δ₀, weak ligands ↓ Δ₀.
• Geometry affects splitting: tetrahedral complexes have smaller splitting
  Δₜ = 4/9 Δ₀ (e.g., Co²⁺: pink → blue change).

CFT is an electrostatic model that considers the metal-ligand bond to be ionic, arising purely from electrostatic interactions between the metal ion and the ligand (treated as point charges for anions, or point dipoles for neutral molecules).

CFT considers the effect of ligands on the relative energies of the d-orbitals of the central metal atom/ion.

If Δ₀ < P, 4th electron will enter e_g giving the configuration \[t_{2g}^3e_{g}^1.\] Ligands for which Δ₀ < P are called weak field ligands.

If Δ₀ > P, pairing will occur in the t_2g orbitals and e_g orbitals will remain vacant. So, the configuration for 4th e⁻ will be \[t_{2g}^4e_{g}^0.\]. For Δ₀ > P, ligands are strong field ligands.

Splitting of d-orbitals in a square planar crystal field:

Splitting of d-orbital in a tetrahedral crystal field: