Transition elements Revision Notes

Transition metals generally have high densities compared to other metals. Density is defined as mass per unit volume (density = mass/volume). The high density of transition metals is due to their closely packed crystal structures and the presence of many electrons, including d-electrons, which contribute to strong metallic bonding.

For example, iron has a density of approximately 7.9 g/cm³, while copper is about 8.9 g/cm³, both significantly denser than metals like sodium (0.97 g/cm³).

The strong metallic bonding arises from the delocalisation of both the outer s-electrons and the d-electrons, resulting in a greater attraction between atoms and thus a more compact structure.

• Density = mass/volume, typically high for transition metals.
• Strong metallic bonding due to delocalised d-electrons.
• Close-packed crystal lattices contribute to high density.

Transition metals have relatively high melting points compared to other metals. This is because the metallic bonding in transition metals is stronger, involving delocalised electrons from both the outer s and d orbitals.

Strong metallic bonds require more energy to break, resulting in higher melting points. For example, tungsten has a melting point of 3422 °C, one of the highest among metals, while iron melts at 1538 °C.

The melting point depends on the strength of the metallic bonds and the crystal structure. The presence of partially filled d orbitals allows more electrons to participate in bonding, increasing bond strength.

• Melting points are high due to strong metallic bonding.
• Delocalised d-electrons increase bond strength.
• Examples: Tungsten (3422 °C), Iron (1538 °C).

Transition metals form coloured compounds because their ions have partially filled d orbitals. When light passes through or reflects off these compounds, electrons in the d orbitals absorb specific wavelengths of visible light to move between different d energy levels (d–d electronic transitions).

The absorbed wavelengths correspond to the energy difference between split d orbitals caused by the ligand field (crystal field splitting). The remaining transmitted or reflected light is the complementary colour, which we observe.

For example, the hydrated copper(II) ion [Cu(H₂O)₆]²⁺ appears blue because it absorbs light in the red/orange region. Similarly, chromium(III) compounds often appear green or violet depending on the ligands and oxidation state.

• Partially filled d orbitals allow d–d electronic transitions.
• Ligand field splitting causes energy differences between d orbitals.
• Selective absorption of visible light leads to observed colours.
• Example: [Cu(H₂O)₆]²⁺ is blue due to absorption of red/orange light.

Transition metals and their compounds often act as catalysts, increasing the rate of chemical reactions without being consumed. This catalytic ability arises from their variable oxidation states and ability to adsorb reactant molecules onto their surfaces.

In their elemental form, transition metals provide active sites where reactants can adsorb, weakening bonds and lowering activation energy. For example, finely divided platinum catalyses the hydrogenation of alkenes.

Transition metal compounds can also act as homogeneous catalysts. For instance, iron(III) ions catalyse the decomposition of hydrogen peroxide: 2 H₂O₂ (aq) → 2 H₂O (l) + O₂ (g). The catalyst cycles between oxidation states during the reaction, facilitating electron transfer processes.

• Catalysts speed up reactions without being consumed.
• Elemental metals adsorb reactants, weakening bonds.
• Variable oxidation states enable redox catalysis in compounds.
• Example: Platinum catalyses alkene hydrogenation.
• Example: Fe³⁺ catalyses H₂O₂ decomposition.

Transition metals have partially filled d-orbitals, allowing them to lose different numbers of electrons when forming ions. This results in multiple stable oxidation states for the same element. The ability to exhibit variable oxidation states is a defining characteristic of transition metals and influences their chemical behaviour, including the formation of coloured compounds and catalytic activity.

Iron, a typical transition metal, commonly shows two oxidation states: +2 and +3. These states arise from the loss of electrons from the 4s and 3d orbitals during ion formation.

• Transition metals have variable oxidation states due to d-electrons.
• Iron commonly exhibits +2 and +3 oxidation states.
• Oxidation states affect chemical properties and compound formation.

The ferrous ion (Fe²⁺) forms when iron loses two electrons, typically from the 4s orbital and one 3d electron. This ion is stable in reducing environments and forms compounds such as iron(II) sulfate (FeSO₄) and iron(II) chloride (FeCl₂).

Iron(II) compounds are usually pale green or blue-green in colour. For example, iron(II) sulfate crystals are pale green.

Fe²⁺ ions can be oxidised to Fe³⁺ ions by oxidising agents such as oxygen or potassium manganate(VII) in acidic solution.

• Fe → Fe²⁺ + 2e⁻ (loss of two electrons)
• Fe²⁺ compounds are typically pale green or blue-green.
• Fe²⁺ is stable in reducing conditions but can be oxidised to Fe³⁺.

The ferric ion (Fe³⁺) forms when iron loses three electrons, including two from the 3d orbitals and one from the 4s orbital. This ion is more stable in oxidising environments.

Iron(III) compounds, such as iron(III) chloride (FeCl₃) and iron(III) oxide (Fe₂O₃), are often yellow, brown or reddish in colour. For example, rust is mainly iron(III) oxide, which is reddish-brown.

Fe³⁺ ions can be reduced back to Fe²⁺ ions by reducing agents such as zinc or iodide ions.

• Fe → Fe³⁺ + 3e⁻ (loss of three electrons)
• Fe³⁺ compounds are typically yellow, brown or reddish.
• Fe³⁺ is stable in oxidising conditions and can be reduced to Fe²⁺.

Iron ions participate in redox reactions where they change oxidation states. For example, iron(II) ions can be oxidised to iron(III) ions by potassium manganate(VII) in acidic solution:

The balanced ionic equation is: 5Fe²⁺(aq) + MnO₄⁻(aq) + 8H⁺(aq) → 5Fe³⁺(aq) + Mn²⁺(aq) + 4H₂O(l)

Conversely, iron(III) ions can be reduced to iron(II) ions by zinc metal: Fe³⁺(aq) + e⁻ → Fe²⁺(aq) (half-equation) and Zn(s) → Zn²⁺(aq) + 2e⁻ (half-equation). The overall reaction is: 2Fe³⁺(aq) + Zn(s) → 2Fe²⁺(aq) + Zn²⁺(aq).

• Oxidation: Fe²⁺ → Fe³⁺ + e⁻
• Reduction: Fe³⁺ + e⁻ → Fe²⁺
• Example redox reaction with MnO₄⁻ in acid: 5Fe²⁺ + MnO₄⁻ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O
• Reduction by zinc: 2Fe³⁺ + Zn → 2Fe²⁺ + Zn²⁺

The different oxidation states of iron produce compounds with distinct colours due to differences in electron configuration and ligand field effects. Fe²⁺ compounds are generally pale green or blue-green, while Fe³⁺ compounds are yellow, brown or reddish.

The reactivity of iron ions depends on their oxidation state. Fe²⁺ ions are more easily oxidised, while Fe³⁺ ions are more easily reduced. This redox flexibility is important in biological systems and industrial processes.

The variable oxidation states also influence the types of complexes iron forms, affecting their stability and function.

• Fe²⁺ compounds: pale green/blue-green colour
• Fe³⁺ compounds: yellow, brown or reddish colour
• Fe²⁺ is more reducing; Fe³⁺ is more oxidising
• Oxidation state affects complex formation and stability