Aluminium and amphoteric behaviour

Aluminium sits on the metal / non-metal borderline, so Al₂O₃ and Al(OH)₃ are amphoteric — they react with both acids and alkalis.

Aluminium metal resists corrosion because a thin, tough, impermeable layer of Al₂O₃ forms instantly on the surface — anodising thickens it deliberately.

Group 4 and the inert pair effect: going down C → Si → Ge → Sn → Pb, the +2 oxidation state becomes progressively more stable relative to +4. So Sn²⁺ is a reducing agent (it wants to become Sn⁴⁺), whereas Pb⁴⁺ is an oxidising agent (it wants to become Pb²⁺).

What does amphoteric mean?

Oxides of Period 3

Across Period 3 the oxides change from ionic and basic to covalent and acidic:

• Na₂O and MgO — ionic, basic. Na₂O dissolves to give a strongly alkaline solution (pH ≈ 13); MgO is only sparingly soluble (pH ≈ 9).
• Al₂O₃ — amphoteric, and insoluble in water.
• SiO₂ — giant covalent, weakly acidic, insoluble in water but attacked by hot concentrated alkali.
• P₄O₁₀ and SO₃ — simple covalent, strongly acidic; they react vigorously with water to give phosphoric(V) acid and sulfuric acid (pH ≈ 0–2).

The chlorides follow the same story: NaCl and MgCl₂ simply dissolve, but AlCl₃, SiCl₄ and PCl₅ are hydrolysed, fuming in moist air and producing acidic solutions of HCl.

Acidic, basic or amphoteric?

Tap an oxide, then the box that describes its reaction with water, acid and alkali.

What makes a transition metal?

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

That single definition explains all four characteristic properties:

• Variable oxidation states — the 4s and 3d orbitals are very close in energy, so a variable number of electrons can be lost.
• Coloured ions — d–d transitions are only possible if the d subshell is partly filled.
• Catalytic activity — variable oxidation states let them accept and donate electrons; and their surfaces adsorb reactants.
• Complex ion formation — the empty d orbitals accept lone pairs from ligands.

The two exceptions: Sc forms only Sc³⁺, which is 3d⁰ (empty), and Zn forms only Zn²⁺, which is 3d¹⁰ (full). Neither has a partly filled d subshell in its ion, so neither is a transition metal — and their compounds are white.

Why is zinc not a transition metal?

Complex ions, ligands and shapes

A ligand is a species with a lone pair that forms a dative covalent (coordinate) bond to a central metal ion. The coordination number is the number of coordinate bonds.

• Monodentate — one donor atom: H₂O, NH₃, Cl⁻, CN⁻, OH⁻.
• Bidentate — two donor atoms: ethane-1,2-diamine ("en"), the ethanedioate ion.
• Multidentate — EDTA⁴⁻ is hexadentate; haem is tetradentate.

Shapes: 6 coordinate bonds → octahedral (90°), the commonest, with small ligands like H₂O and NH₃. 4 coordinate bonds → usually tetrahedral (109.5°) with the bulkier Cl⁻; but Pt(II) complexes such as cisplatin are square planar. 2 → linear, e.g. [Ag(NH₃)₂]⁺ in Tollens' reagent.

The chelate effect: a multidentate ligand displaces monodentate ones because the reaction produces more particles — a large positive entropy change makes ΔG negative.

Count the bonds

Your turn — oxidation state in a complex

Why transition metal complexes are coloured

Ligands split the five d orbitals into two groups of slightly different energy. An electron can absorb a photon of visible light and jump the gap (a d–d transition) — but only if the d subshell is partly filled, so there is both an electron to promote and a space to promote it into.

The colour you see is the complementary colour to the light absorbed. Change the ligand, the oxidation state or the coordination number and you change ΔE — and therefore the colour.

Colorimetry: because the absorbance is proportional to concentration, a colorimeter (with a filter of the complementary colour) can measure the concentration of a coloured complex.

Your turn — energy of a d-d transition

Ligand substitution

One ligand replaces another. If the new ligand is a similar size (H₂O and NH₃), the coordination number is unchanged; if it is much larger (Cl⁻), the coordination number usually drops from 6 to 4.

Carbon monoxide and haemoglobin: the Fe(II) in haem normally binds O₂ reversibly. CO binds to the same site far more strongly, so it is not readily displaced — the blood can no longer carry oxygen. This is a ligand substitution reaction.

Add concentrated HCl

Transition metals as catalysts

• Heterogeneous — the catalyst is in a different phase from the reactants. Reactants adsorb onto active sites on the surface, bonds weaken, they react, and the products desorb. Examples: Fe in the Haber process, V₂O₅ in the Contact process, Ni in hydrogenation, Pt/Rh/Pd in a catalytic converter. Increase the surface area to increase activity — and beware poisoning (sulfur poisons the Haber iron; lead poisons a catalytic converter).
• Homogeneous — the catalyst is in the same phase, and works by forming an intermediate. Example: Fe²⁺ catalysing the reaction between S₂O₈²⁻ and I⁻ (two low-activation-energy steps replace one high-E_a collision between two negative ions).
• Autocatalysis — a product catalyses the reaction. Mn²⁺ autocatalyses the MnO₄⁻ / ethanedioate reaction, so the rate speeds up as the reaction proceeds.

All of this depends on variable oxidation states — the metal can accept electrons and then give them back, returning unchanged at the end.

Name the catalyst

Redox chemistry of the transition metals

Variable oxidation states make transition metals ideal for redox titrations.

Dichromate titrations need an indicator (e.g. barium diphenylamine sulfonate) because green Cr³⁺ masks the colour change.

Your turn — dichromate titration

Your turn — percentage of iron

Match the colour

Tap an ion or complex on the left, then the colour you would see.

The big ideas to know

That is the whole of PI2. Press Finish to see your score.