Properties of Transition Elements HL Only

Full Notes

Definition

A transition element is defined as a d-block element that forms at least one ion with a partially filled d sublevel.

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.

Key Properties Explained

Variable Oxidation States

Due to similar energies of 4s and 3d sublevels, transition metals can lose different numbers of d and s electrons.

Examples:

• Fe: Fe²⁺ (+2), Fe³⁺ (+3)
• Mn: up to +7 in MnO₄⁻

High Melting Points

Strong metallic bonding involving delocalized d-electrons results in high melting points.

Magnetic Properties

Magnetism arises when unpaired electrons generate tiny magnetic fields as they spin, which can interact with external magnetic fields.

In transition metals, this effect comes from unpaired d-electrons. The more unpaired electrons, the stronger the magnetic effect.

You don’t need to know the names of specific types of magnetism for IB exams, but you should understand that magnetism is closely linked to electron configuration, especially the number of unpaired d-electrons.

Catalytic Properties

Catalysts increase the rate of a reaction by providing an alternative reaction pathway with a lower activation energy (E_a).

Transition metals are often used as catalysts because of their ability to form ions with different oxidation states and because of their relatively low reactivity.

Examples:

• Fe in Haber process
• V₂O₅ in the Contact process

Colour in Transition Metal Complexes

This is covered in more detail in S3.1.10

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 orbital shapes experience differing amounts of repulsion, meaning the orbitals split into different energies. There is an energy gap (ΔE) between the d-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 happen 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 transitions can occur

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 its lone pairs 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⁺

A central Cu²⁺ ion is surrounded by 6 water ligands.

The formulas of complex ions are written in square brackets with the overall charge of the complex ion shown as a superscript.

Summary Table of Properties

Summary

• Transition elements are defined by partially filled d sublevels in at least one ion
• They show variable oxidation states and high melting points
• Magnetic properties come from unpaired d-electrons
• They act as catalysts due to oxidation state flexibility
• They form coloured compounds via d-d transitions
• They form complex ions with ligands through coordinate bonds