Bonding in Coordination Compounds

Full Notes

Valence Bond Theory

This theory explains the bonding in coordination compounds by considering hybridisation of atomic orbitals in the metal atom or ion. Ligands donate electron pairs into these hybrid orbitals to form coordinate bonds, resulting in specific geometries.

• The central metal ion makes a number of vacant orbitals available for bonding.
• Hybridisation of these orbitals occurs to form equivalent hybrid orbitals.
• Ligands donate lone pairs into these hybrid orbitals, forming coordinate covalent bonds.
• The geometry of the complex is determined by the type of hybridisation:
  • sp³ = tetrahedral
  • dsp² = square planar
  • d²sp³ = inner orbital octahedral
  • sp³d² = outer orbital octahedral

Examples:

[Co(NH₃)₆]³⁺

• Co³⁺: 3d⁶ configuration
• Hybridisation: d²sp³ (inner orbital)
• Geometry: Octahedral
• Magnetic behaviour: Diamagnetic (no unpaired electrons)

[CoF₆]³⁻

• Co³⁺: 3d⁶ configuration
• F⁻ is a weak field ligand → no pairing
• Hybridisation: sp³d² (outer orbital)
• Geometry: Octahedral
• Magnetic behaviour: Paramagnetic (4 unpaired electrons)

[NiCl₄]²⁻

• Ni²⁺: 3d⁸ electronic configuration
• Cl⁻ is a weak field ligand → no pairing of electrons
• Hybridisation: sp³ (outer orbital hybridisation)
• Geometry: Tetrahedral
• Magnetic behaviour: Paramagnetic (with 2 unpaired electrons)

[Ni(CN)₄]²⁻

• Ni²⁺: 3d⁸ electronic configuration
• CN⁻ is a strong field ligand → pairing occurs
• Hybridisation: dsp² (inner orbital hybridisation)
• Geometry: Square planar
• Magnetic behaviour: Diamagnetic (no unpaired electrons)

Magnetic Properties of Coordination Compounds

The magnetic behaviour of a coordination compound depends on the presence or absence of unpaired electrons in its d-orbitals. This can be measured experimentally and helps in predicting the electronic configuration and nature of the complex.

d-Electron Configurations and Hybridisation:

• For metal ions with up to 3 electrons in d orbitals (e.g. Ti³⁺, V³⁺, Cr³⁺ → d¹, d², d³):
  • Two vacant d-orbitals are available for octahedral hybridisation (with 4s and 4p orbitals).
  • Their magnetic behaviour is consistent whether free or in complexes.
• When more than 3 d-electrons are present:
  • Hund’s rule restricts pairing, making inner d-orbital hybridisation less accessible.
  • For d⁴ (e.g. Cr²⁺, Mn³⁺), d⁵ (e.g. Mn²⁺, Fe³⁺), d⁶ (e.g. Fe²⁺, Co³⁺): pairing for inner orbital hybridisation is difficult, leading to outer orbital hybridisation instead.

Limitations of Valence Bond Theory

Although valence bond theory (VBT) explains bonding and magnetic properties to some extent, it has several limitations that restrict its use with more advanced coordination chemistry.

• It does not explain the colour of complexes.
• It provides no insight into the relative strengths of ligands.
• It does not distinguish between high-spin and low-spin complexes.
• It cannot predict the spectra of complexes.
• It ignores the quantitative aspects of bonding and thermodynamic stability.

Crystal Field Theory

Crystal Field Theory models metal-ligand bonding as purely electrostatic, explaining how the presence of ligands causes splitting of d-orbitals. It provides a better understanding of the geometry, colour, and magnetism of complexes.

• Ligands are treated as point charges (if anionic) or dipoles (if neutral molecules).
• In a spherical field of negative charge, all five degenerate d-orbitals of a free metal ion experience equal repulsion from the surrounding negative field, resulting in a uniform increase in their energy.
• However, in the presence of directional fields—such as octahedral, tetrahedral, or square planar geometries—the d-orbitals experience varying degrees of repulsion depending on their spatial orientation.
• Because each d-orbital is shaped differently, some interact more strongly with the ligand field than others. This causes the previously degenerate (equal-energy) d-orbitals to split into two distinct sets: one with higher energy and one with lower energy.
• When ligands coordinate to a central metal ion, this crystal field effect causes the d-orbitals to split based on the geometry and strength of the ligand field.

The extent of crystal field splitting (Δ_o) depends on two main factors:

1. The nature of the ligand (i.e., how strong or weak its field is), and
2. The charge on the metal ion (higher charge → greater splitting).

Some ligands produce a strong field, causing a large energy gap between the split d-orbitals. Others produce a weaker field, leading to smaller splitting.

The spectrochemical series is an experimentally determined ranking of ligands from weakest to strongest field strength based on how much they split the d-orbitals in metal complexes.

Example:
I⁻ < Br⁻ < SCN⁻ < Cl⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < CN⁻ < CO

Octahedral Field Splitting

Five d-orbitals split into:

• t_2g (d_xy, d_yz, d_xz) = lower energy
• e_g (d_z2, d_x2−y2) = higher energy

• The energy difference is called crystal field splitting energy (Δ_o).
• Strong field ligands (e.g., CN⁻) give large Δ_o and pairing of electrons in orbitals occurs = low-spin complex
• Weak field ligands (e.g., F⁻) give small Δ_o and pairing of electrons in orbitals does not occur = high-spin complex

Tetrahedral Field Splitting

Reverse splitting:

• t₂ (d_x2−y2, d_z2) = lower energy
• e (d_xy, d_yz, d_xz) = higher energy

• Crystal field splitting (Δ_t) is smaller than Δ_o, so high-spin complexes are favoured.

Crystal Field Stabilisation Energy (CFSE): Difference in energy between the d-orbital configuration in free ion vs complex. Explains stability and colour trends.

Colour in Coordination Compounds

The colours of coordination compounds are due to electronic transitions within the d-orbitals.

When visible light is absorbed to promote electrons between split d-levels, the complementary colour is observed.

• Most coordination compounds are coloured due to d–d transitions.
• When white light strikes a complex:
  • One wavelength is absorbed (corresponding to electronic transition between d-orbitals)
  • The complementary colour is observed
• Example: [Ti(H₂O)₆]³⁺ appears purple due to absorption in the green-yellow region.

Note: If all d-electrons are paired, or if no d-electrons are present, the complex is colourless.

Limitations of Crystal Field Theory

• It assumes purely electrostatic interaction, ignoring covalent character in bonding.
• It cannot explain the spectral and thermodynamic properties of all complexes.
• It overestimates the ionic nature of metal-ligand bonds.
• Fails to explain the ligand field strength order (spectrochemical series) in terms of bonding.

Summary

• Valence Bond Theory describes hybridisation and predicts geometry and magnetism.
• Crystal Field Theory explains d-orbital splitting, colour and spin states.
• Strong field ligands give large splitting and often low-spin complexes.
• Weak field ligands give small splitting and often high-spin complexes.
• Both VBT and CFT have limitations for stability and reactivity trends.