Transition Metals

This lesson covers the properties of d-block elements, including variable oxidation states, coloured ions, complex ion formation, ligand substitution, catalysis, isomerism in complex ions, precipitation reactions for identifying ions, redox titrations, and colourimetry. The transition metals exhibit unique chemistry due to the involvement of their partially filled d orbitals. This material aligns with the AQA and OCR A specifications for A-Level Chemistry.

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What Is a Transition Metal?

Key Definition: A transition metal is a d-block element that forms at least one stable ion with a partially filled d sub-shell.

Note that scandium (Sc) and zinc (Zn) are d-block elements but are not transition metals by this definition — Sc³⁺ has an empty d sub-shell (3d⁰) and Zn²⁺ has a full d sub-shell (3d¹⁰). Copper is a transition metal because, although Cu⁺ has 3d¹⁰ (full), Cu²⁺ has 3d⁹ (partially filled).

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Properties of Transition Metals

Transition metals share several characteristic properties, all arising from the involvement of the partially filled 3d sub-shell:

1. Variable oxidation states — transition metals can form ions with different charges because electrons can be lost from both the 4s and 3d sub-shells. The common oxidation states of the first-row transition metals include:

Element	Common oxidation states	Most stable
Ti	+3, +4	+4
V	+2, +3, +4, +5	+5
Cr	+2, +3, +6	+3
Mn	+2, +3, +4, +6, +7	+2
Fe	+2, +3	+3
Co	+2, +3	+2
Ni	+2	+2
Cu	+1, +2	+2

2. Formation of coloured ions — when ligands surround a transition metal ion, the five d orbitals (which are normally degenerate, i.e. the same energy) split into two groups of different energies. Electrons can absorb visible light and be promoted from the lower-energy d orbitals to the higher-energy ones. This is called a d–d transition. The colour observed is the complementary colour of the wavelength of light absorbed.

Diagram description (d-orbital splitting in an octahedral complex): The five d orbitals split into two groups: three lower-energy orbitals (labelled t₂g: dxy, dxz, dyz) and two higher-energy orbitals (labelled eg: dx²−y², dz²). The energy gap between them is labelled ΔE (or Δoct). When ΔE corresponds to the energy of visible light, the complex is coloured.

3. Catalytic activity — transition metals and their compounds are effective catalysts because they can change oxidation state (providing an alternative reaction pathway with lower activation energy) and because they can adsorb reactants onto their surface (in heterogeneous catalysis), bringing them into close proximity and weakening their bonds. Examples:

   • Iron in the Haber process (N₂ + 3H₂ ⇌ 2NH₃)
   • Vanadium(V) oxide (V₂O₅) in the Contact process (2SO₂ + O₂ ⇌ 2SO₃)
   • Manganese(IV) oxide (MnO₂) in the decomposition of hydrogen peroxide
   • Nickel in the hydrogenation of alkenes and vegetable oils

4. Formation of complex ions — see below.

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Complex Ions

Key Definition: 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 Ligand

Type	Definition	Examples
Monodentate	Donates one lone pair per ligand	H₂O, NH₃, Cl⁻, CN⁻, CO, OH⁻
Bidentate	Donates two lone pairs per ligand	Ethane-1,2-diamine (en, H₂NCH₂CH₂NH₂), ethanedioate (ox, C₂O₄²⁻)
Multidentate (polydentate)	Donates multiple lone pairs per ligand	EDTA⁴⁻ (hexadentate — donates 6 lone pairs)

Coordination Number

The coordination number is the total number of dative bonds from ligands to the central metal ion. Common coordination numbers and their corresponding shapes:

Coordination number	Shape	Bond angle	Example
2	Linear	180°	[Ag(NH₃)₂]⁺
4	Tetrahedral	109.5°	[CoCl₄]²⁻, [CuCl₄]²⁻
4	Square planar	90°	[Pt(NH₃)₂Cl₂] (cisplatin)
6	Octahedral	90°	[Cu(H₂O)₆]²⁺, [Fe(CN)₆]⁴⁻, [Cr(H₂O)₆]³⁺

Isomerism in Complex Ions: Cis-Trans Isomerism

Square planar and octahedral complexes with two or more different types of ligand can show cis-trans isomerism (a type of geometric/stereoisomerism).

Cis-trans isomerism in square planar complexes:

Consider [Pt(NH₃)₂Cl₂]:

• Cis isomer: the two NH₃ ligands (and the two Cl⁻ ligands) are on the same side (adjacent, 90° apart).
• Trans isomer: the two NH₃ ligands (and the two Cl⁻ ligands) are on opposite sides (180° apart).

Diagram description: Draw a square with Pt at the centre. In the cis form, the two Cl ligands occupy adjacent corners (e.g. top and right), and the two NH₃ ligands occupy the other two adjacent corners. In the trans form, the two Cl ligands are at opposite corners (e.g. top and bottom), and the two NH₃ are at the other opposite corners.

Biological significance of cisplatin: The cis isomer of [Pt(NH₃)₂Cl₂] is the anti-cancer drug cisplatin. It works by binding to DNA strands in cancer cells — the two Cl⁻ ligands are replaced by nitrogen atoms in guanine bases on the same DNA strand, cross-linking the strand and preventing DNA replication. The trans isomer does not have the correct geometry to bind to DNA in this way and is therefore not an effective anti-cancer agent.

Cis-trans isomerism in octahedral complexes also occurs — for example, [Co(NH₃)₄Cl₂]⁺ can exist as cis and trans forms, where the two Cl⁻ ligands are either adjacent (90°) or opposite (180°).

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Ligand Substitution

Ligands can be exchanged (substituted) in complex ions. This can cause a change in colour, coordination number, and shape.

Common Ligand Substitution Reactions

1. Copper complexes:

[Cu(H₂O)₆]²⁺ (pale blue) + 4Cl⁻ ⇌ [CuCl₄]²⁻ (yellow-green) + 6H₂O

Coordination number changes from 6 to 4; shape changes from octahedral to tetrahedral (because Cl⁻ is larger than H₂O, so fewer Cl⁻ ligands can fit around the Cu²⁺ ion).

[Cu(H₂O)₆]²⁺ (pale blue) + 4NH₃ → [Cu(NH₃)₄(H₂O)₂]²⁺ (deep blue/violet) + 4H₂O

Here only 4 of the 6 water molecules are replaced by NH₃ (partial substitution); the coordination number remains 6 and the shape remains octahedral.

graph TD
    A["[Cu(H₂O)₆]²⁺<br/>Pale blue<br/>Octahedral, CN=6"] -->|"+ 4Cl⁻"| B["[CuCl₄]²⁻<br/>Yellow-green<br/>Tetrahedral, CN=4"]
    A -->|"+ 4NH₃"| C["[Cu(NH₃)₄(H₂O)₂]²⁺<br/>Deep blue<br/>Octahedral, CN=6"]
    A -->|"+ excess NaOH"| D["Cu(OH)₂ precipitate<br/>Blue<br/>(insoluble in excess)"]

2. Cobalt complexes:

[Co(H₂O)₆]²⁺ (pink) + 4Cl⁻ ⇌ [CoCl₄]²⁻ (blue) + 6H₂O

3. Chromium complexes:

[Cr(H₂O)₆]³⁺ (green/violet) + 6NH₃ → [Cr(NH₃)₆]³⁺ (purple) + 6H₂O

The Chelate Effect

Key Definition: The chelate effect is the increased stability of complexes formed with bidentate or multidentate ligands compared to those formed with monodentate ligands.

The increased stability is due to the entropy increase when one multidentate ligand replaces several monodentate ligands — the total number of free species in solution increases.

For example, when three bidentate ethane-1,2-diamine (en) ligands replace six water molecules:

[Ni(H₂O)₆]²⁺ + 3en → [Ni(en)₃]²⁺ + 6H₂O

This goes from 4 species (1 complex + 3 en) to 7 species (1 complex + 6 H₂O). The increase in the number of particles gives a positive ΔS, making ΔG more negative and the equilibrium constant larger. The chelate complex is therefore more thermodynamically stable.

When EDTA⁴⁻ (hexadentate) replaces six water molecules: 2 species → 7 species. The entropy increase is even greater.

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