Principles of Transition Metal Chemistry
Quick Notes
- Transition metals are d-block elements that form stable ions with incomplete d orbitals.
- Their electronic configurations follow the pattern where 4s fills before 3d, but 4s is lost first when ions form.
- They show variable oxidation states due to similar energies of orbitals in their 4s and 3d sub-shells.
- A ligand is a molecule or ion that donates a lone pair to a central metal ion.
- Coordinate (dative covalent) bonding forms between a ligand and the metal ion in a complex ion.
- Transition metal complexes are often coloured due to electron transitions between split d-orbitals.
- Colour changes occur with changes in oxidation number, ligand type, or coordination number.
- Monodentate ligands donate one pair (e.g. H2O, NH3); bidentate donate two (e.g. en); multidentate like EDTA4− donate more.
- Coordination number = number of ligand bonds to the metal. Common values: 4 or 6.
- Octahedral, tetrahedral, and square planar are typical complex geometries.
- Cis-platin, a square planar complex, is used as a chemotherapy drug — the cis isomer is active.
- Haemoglobin is an Fe2+ complex with a multidentate ligand. It binds O2 reversibly, but CO binds irreversibly, causing toxicity.
Full Notes
Electronic Configurations of d-Block Elements
The 4s orbital fills before the 3d orbital when building up atoms, but 4s electrons are lost first when ions form.
Example Fe and Fe2+
Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d6
Fe2+: 1s2 2s2 2p6 3s2 3p6 3d6
Definition of a Transition Metal
Transition metals are d-block elements that form stable ions with partially-filled d-orbitals.
Scandium and zinc do not meet this definition in all oxidation states:
- Sc3+: 3d0 (no electrons in d-orbital)
- Zn2+: 3d10 (full d-orbital)
Hence, they are not considered transition metals in their common ions.
Variable Oxidation States
The small energy gap between 4s and 3d orbitals means different numbers of electrons can be readily lost, leading to multiple oxidation states.
E.g., manganese (Mn) can form a range of oxidation states from +2 to +7.
This makes transition metals useful in redox reactions and catalysis.
Ligands and Complex Ion Formation
A ligand is a molecule or ion that donates a lone pair to form a coordinate bond with a metal ion.
For example, water molecules (H2O) are able to act as ligands as the oxygen atom can use one of its lone pairs of electrons to form a co-ordinate bond to a central metal atom or ion.
Complex ions are formed when a metal ion is surrounded by ligands via coordinate (dative covalent) bonds.
Example [Cu(H2O)6]2+
![CIE A-Level Chemistry octahedral complex [Cu(H2O)6]2+ with six aqua ligands around Cu2+.](images/coppercomplex.png)
This is a copper ion surrounded by six water ligands.
The formulas of complex ions are written in square brackets with the overall charge of the complex ion shown as a superscript.
Colour in Transition Metal Complexes
When ligands bond to a metal ion, the ion’s d-orbitals split into two energy levels (higher and lower).
This occurs because electrons in the d-orbitals are repelled by electrons from incoming ligands.
- Different d-orbital shapes experience differing amounts of repulsion, causing an energy gap (ΔE) to form between the orbitals.
Electrons can absorb energy from visible light to move from a lower energy level (ground state) to a higher one (excited state).
The remaining wavelengths of light are transmitted or reflected, giving the solution its observed colour.
Colour changes occur when:
- The oxidation state of the metal changes (e.g. Fe2+ vs Fe3+).
- The ligand changes (e.g. H2O vs NH3).
- The coordination number changes (e.g. 6 → 4 ligands).
No colour is seen if the metal has a full (d10) or empty (d0) d sub-shell, so no electron transitions can occur.
Coordination Number and Complex Shapes
Co-ordination number refers to the number of co-ordinate bonds around a central metal ion and determines the geometry (shape) of the complex.
Shapes depend on ligand size and number:
- Octahedral (6 ligands), common with small ligands (e.g. H2O, NH3)
- Tetrahedral (4 ligands), often with larger ligands like Cl−
- Square planar (4 ligands), typical for Pt2+, e.g. cis-platin
Most complexes have a co-ordination number of 6 (octahedral) or 4 (tetrahedral or square planar).
Don’t confuse co-ordination number with the number of ligands in a complex ion! Sometimes the number of ligands can be different to the co-ordination number (for complexes with bidentate and multidentate ligands in).
Cis-Platin as an Anti-Cancer Drug
Cis-platin ([Pt(NH3)2Cl2]) has Cl− ligands on the same side (cis isomer).
It binds to DNA in cancer cells, blocking replication.
Trans-platin is not effective, as it doesn't bind DNA in the same way.
Types of Ligand
We classify ligands as monodentate, bidentate or multidentate, based on the number of co-ordinate bonds they can form.
Monodentate Ligands
Monodentate ligands can donate one lone pair to the central metal ion.
Common examples:
Bidentate Ligands
Bidentate ligands form two co-ordinate bonds, with two atoms donating one lone pair of electrons.
Common examples: Ethan-1,2-diamine (“en") and ethanedioate
Multidentate Ligands
Multidentate ligands form multiple co-ordinate bonds.
Common Example: EDTA4− (forms 6 co-ordinate bonds)
Haemoglobin and Oxygen Transport
Haemoglobin is a multidentate Fe2+ complex found in red blood cells that binds to oxygen via reversible ligand exchange.
The oxygen (O2) binds to Fe2+ via a co-ordinate bond, allowing oxygen transport in the blood.
CO is toxic because it binds more strongly than O2.
Summary
- Transition metals are d-block elements that form ions with incomplete d orbitals and show variable oxidation states.
- Ligands donate lone pairs to form coordinate bonds, creating coloured complexes whose shapes depend on coordination number.
- Changing oxidation state, ligand or coordination number can change a complex’s colour.
- Cis-platin is an important anti-cancer drug and haemoglobin is a biological Fe2+ complex that transports oxygen but binds CO too strongly.