Transition Metals and Complex Ions

A transition metal is a d-block element that forms at least one stable ion with a partially filled d subshell.

• Zinc is NOT a transition metal: Zn²⁺ = $[\text{Ar}] 3d^{10}$ (full d shell)
• Scandium is debatable: Sc³⁺ = $[\text{Ar}]$ (empty d shell)
• Copper IS a transition metal: Cu²⁺ = $[\text{Ar}] 3d^9$ (partially filled)

A complex ion consists of a central metal ion surrounded by ligands.

A ligand is a molecule or ion that donates a lone pair of electrons to the metal ion, forming a dative (coordinate) bond.

Types of Ligands

• Monodentate: donates 1 lone pair
• Bidentate: donates 2 lone pairs
• Hexadentate (polydentate): donates 6 lone pairs

Coordination Number

The coordination number is the total number of dative bonds from ligands to the central metal ion.

• 6 coordinate → octahedral shape (e.g. $[\text{Cu(H}_2\text{O)}_6]^{2+}$)
• 4 coordinate → tetrahedral (e.g. $[\text{CoCl}_4]^{2-}$) or square planar (e.g. cisplatin)
• 2 coordinate → linear (e.g. $[\text{Ag(NH}_3\text{)}_2]^+$)

Why Are Transition Metal Compounds Coloured?

In a complex ion, the d orbitals split into two energy levels. When visible light hits the complex:

1. An electron in a lower d orbital absorbs energy
2. It jumps to a higher d orbital (d-d transition)
3. The wavelength absorbed depends on the energy gap ($\Delta$)
4. The complementary colour is what we see

Factors Affecting Colour

• Oxidation state of the metal ion (changes d electron configuration)
• Ligand (different ligands cause different d-orbital splitting)
• Coordination number (affects geometry and splitting)

Ligands can be replaced by other ligands. This often causes a colour change.

Example: Copper Complexes

$[\text{Cu(H}_2\text{O)}_6]^{2+} + 4\text{NH}_3 \rightarrow [\text{Cu(NH}_3\text{)}_4\text{(H}_2\text{O)}_2]^{2+} + 4\text{H}_2\text{O}$

Blue → deep blue (only 4 of 6 water molecules replaced)

$[\text{Cu(H}_2\text{O)}_6]^{2+} + 4\text{Cl}^- \rightarrow [\text{CuCl}_4]^{2-} + 6\text{H}_2\text{O}$

Blue → yellow-green (also changes from octahedral to tetrahedral)

Example: Cobalt Complexes

$[\text{Co(H}_2\text{O)}_6]^{2+} + 4\text{Cl}^- \rightarrow [\text{CoCl}_4]^{2-} + 6\text{H}_2\text{O}$

Pink → blue (octahedral → tetrahedral)

Transition metals act as catalysts because they have variable oxidation states, allowing them to accept and donate electrons.

Heterogeneous catalyst: different phase from reactants (usually solid catalyst, gas/liquid reactants) Homogeneous catalyst: same phase as reactants (usually all in solution)

• Know the definition: partially filled d subshell in at least one stable ion
• Learn the colours of common complexes
• Coordination number 6 = octahedral; 4 = tetrahedral or square planar
• Explain colour using d-d transitions and energy gap
• In catalysis questions, emphasise variable oxidation states

• Transition metals have partially filled d orbitals → variable oxidation states, colour, catalytic activity
• Complex ions: central metal + ligands (donate lone pairs via dative bonds)
• Colour from d-d transitions; depends on ligand, oxidation state, and geometry
• Ligand substitution causes colour and geometry changes
• Catalysts: variable oxidation states allow alternative reaction pathways