Bonding in Coordination Compounds — Part 1: Valence bond theory: visual lesson, ISC Class 12 Chemistry

Lesson

How do ligands bond to metals?

Part 1 — valence bond theory: hybrid orbitals, shapes and magnetism of complexes.

In this lesson you will

State the main ideas of valence bond theory for complexes
Predict the hybridisation, shape and magnetic behaviour of a complex
Distinguish inner and outer orbital (low and high spin) complexes
State the limitations of valence bond theory

Concept

Main ideas of valence bond theory

(1) The central metal ion provides a number of empty orbitals equal to its coordination number. (2) These orbitals hybridise to give equivalent orbitals of definite direction. (3) Each ligand donates a lone pair into one hybrid orbital, forming a coordinate (σ) bond. (4) Strong-field ligands (CN⁻, CO, NH₃ with Co³⁺) can force d electrons to pair, freeing inner (n−1)d orbitals; weak-field ligands (F⁻, Cl⁻, H₂O) cannot.

Table of hybridisations and shapes for coordination numbers two, four and six
CN	Hybridisation	Shape
2	sp	Linear
4	sp³	Tetrahedral
4	dsp²	Square planar
6	d²sp³ (inner orbital)	Octahedral
6	sp³d² (outer orbital)	Octahedral

Concept

Inner and outer orbital complexes

For CN 6, if two inner 3d orbitals are used ( $d^2sp^3$ ), the complex is inner orbital (low spin, fewer unpaired electrons). If two outer 4d orbitals are used ( $sp^3d^2$ ), it is outer orbital (high spin). Magnetic moment ( $\mu = \sqrt{n(n+2)}$ BM) tells us which.

\mu = \sqrt{n(n+2)}\ \text{BM}

Table comparing inner and outer orbital octahedral complexes
Feature	Inner orbital	Outer orbital
Hybridisation	d²sp³	sp³d²
d orbitals used	(n−1)d	nd
Ligands	Strong field	Weak field
Unpaired electrons	Fewer (low spin)	More (high spin)
Example	[Fe(CN)₆]⁴⁻	[Fe(H₂O)₆]²⁺

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