Transition Metals

Transition-metal chemistry has a reputation for being one of the hardest parts of A-level chemistry, not because any single idea is difficult, but because it leans on several earlier ones at once. Three are worth refreshing before you start.

General properties

The transition metals are the d-block elements titanium to copper. What sets them apart is a specific electron feature.

You have already met these ions, if only briefly. At GCSE you knew that iron forms Fe²⁺ and Fe³⁺, and that copper(II) sulfate solution is blue. A-level explains what those descriptions were hiding: in water a metal ion is never “bare”. Each one is surrounded by small molecules or ions, so the blue solution you remember is really the ion [Cu(H₂O)₆]²⁺ (a copper ion wrapped in six water molecules) and its colour comes from that arrangement. The surrounding species are called ligands, the whole assembly is a complex, and in the same way “Fe²⁺(aq)” is really [Fe(H₂O)₆]²⁺. Almost everything in this topic follows from that one picture, so the precise definitions come next.

Because the definition needs a partially filled d sub-level in an ion, two d-block elements are not transition metals: scandium (its only ion Sc³⁺ is [Ar], empty d) and zinc (its only ion Zn²⁺ is [Ar]3d¹⁰, full d).

Electron configurations fill 4s before 3d, with two anomalies for extra stability: chromium is [Ar]3d⁵4s¹ and copper is [Ar]3d¹⁰4s¹. On forming ions, the 4s electrons are removed first (e.g. Fe is [Ar]3d⁶4s²; Fe²⁺ is [Ar]3d⁶).

Shapes of complex ions

The shape follows from the co-ordination number, which usually follows from the size of the ligand. Small ligands pack six around the metal; larger ligands fit only four. The shapes themselves come from the same electron-pair repulsion you used for simple molecules at AS: the ligands arrange themselves as far apart as possible around the central ion.

Isomerism. Octahedral complexes of the type MA₄B₂ (and square planar MA₂B₂) show cis–trans isomerism, a special case of E–Z isomerism. Cisplatin is the cis isomer (the two chlorides adjacent); the trans isomer is not an effective drug: the two chlorides have to be next to each other to cross-link two guanine bases in DNA. Octahedral complexes with bidentate ligands (such as [Ni(en)₃]²⁺) show optical isomerism: two non-superimposable mirror images.

Ligand substitution

One ligand can be replaced by another. Whether the co-ordination number changes depends on the size of the incoming ligand.

NH₃ and H₂O are similar in size and both uncharged, so they exchange with no change of co-ordination number (both stay 6). With copper the substitution is incomplete (only four of the six waters are replaced) and the colour deepens (pale blue → deep blue):

[Cu(H₂O)₆]²⁺ + 4NH₃ ⇌ [Cu(NH₃)₄(H₂O)₂]²⁺ + 4H₂O

Cl⁻ is larger, so fewer fit: the co-ordination number changes from 6 to 4 and the colour changes (blue → yellow):

[Cu(H₂O)₆]²⁺ + 4Cl⁻ ⇌ [CuCl₄]²⁻ + 6H₂O

The same 6→4 change happens with cobalt (pink [Co(H₂O)₆]²⁺ → blue [CoCl₄]²⁻) and iron(III) ([Fe(H₂O)₆]³⁺ → [FeCl₄]⁻).

Ligands can donate more than one lone pair: bidentate ligands (ethane-1,2-diamine H₂NCH₂CH₂NH₂, and the ethanedioate ion C₂O₄²⁻) donate two, and multidentate ligands (EDTA⁴⁻) donate several. Haem is an iron(II) complex with a multidentate ligand; oxygen bonds co-ordinately to the Fe(II) in haemoglobin to be carried in the blood, and carbon monoxide is toxic because it replaces that oxygen.

The origin of colour

In a complex, the d orbitals split into two energy levels a small gap apart. When white light passes through, d electrons absorb the wavelengths whose energy matches that gap and are promoted from the ground state to an excited state: a d–d transition. The energy of the gap is:

ΔE = hν = hc/λ

The colour you see is the light that is not absorbed: the remaining wavelengths that are transmitted or reflected. A complex looks yellow because it reflects/transmits yellow light, not because it emits it.

The colours of some common ions:

The absorption of visible light is the basis of spectroscopy. Colorimetry applies it to measure concentration: a ligand is added to give an intense colour, a filter of the complementary colour is chosen, and the absorbance is measured. A calibration graph of absorbance against concentration (from standard solutions) then converts the unknown’s absorbance into its concentration.

Colour changes in ligand substitution

The colour chain now explains the colour changes you met earlier under ligand substitution. Whenever a reaction swaps one ligand for another (or changes the co-ordination number), it alters the d-orbital gap ΔE, so a different wavelength of light is absorbed and the colour shifts.

In each case the reason is the same: a new ligand (and sometimes a new co-ordination number) gives a new ΔE, so a different wavelength is absorbed and a different colour is transmitted. A colour change on adding a reagent therefore signals that the ligand or co-ordination number has changed. A colour change from a change of oxidation state works for the same underlying reason (a new ΔE) and is covered next.

Variable oxidation states & redox

Transition metals form ions in several oxidation states, each usually a different colour. Vanadium is the classic example: acidified vanadate(V) is reduced step by step by zinc, passing through four colours.

Whether one oxidation state can be reduced to another depends on the electrode potentials, and those are influenced by pH and by the ligand. (Tollens’ reagent, [Ag(NH₃)₂]⁺, uses this: it is reduced to silver by an aldehyde but not a ketone.)

Redox titrations

Manganate(VII), MnO₄⁻, is a powerful oxidising agent and is self-indicating: it is decolourised as it reacts, so the end point is the first permanent trace of purple/pink. Two standard titrations:

5Fe²⁺ + MnO₄⁻ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O

2MnO₄⁻ + 5C₂O₄²⁻ + 16H⁺ → 2Mn²⁺ + 10CO₂ + 8H₂O

Run the redox titration below. Because MnO₄⁻ is self-indicating, no indicator is added: every drop is decolourised as it reacts, and the end point is the drop that leaves the first permanent pale pink.

Catalysis

Transition metals and their compounds catalyse many reactions, and their variable oxidation states are the key: the metal can gain and lose electrons, cycling between states as it provides an easier route.

Heterogeneous catalysts

A heterogeneous catalyst is in a different phase from the reactants, and the reaction happens at active sites on its surface (reactants adsorb, react, then the products desorb). A support medium is used to spread the catalyst thinly, maximising the surface area and minimising the cost of the metal. Active sites can be poisoned by impurities that block them, reducing efficiency (a cost implication).

• V₂O₅ catalyses the Contact process (SO₂ → SO₃), working through two steps as vanadium cycles between +5 and +4:

V₂O₅ + SO₂ → V₂O₄ + SO₃

2V₂O₄ + O₂ → 2V₂O₅

• Fe catalyses the Haber process (N₂ + 3H₂ ⇌ 2NH₃).

Homogeneous catalysts

A homogeneous catalyst is in the same phase as the reactants, and the reaction proceeds through an intermediate. The classic example is Fe²⁺ catalysing the reaction between the two negative ions S₂O₈²⁻ and I⁻, which otherwise repel each other and react slowly:

S₂O₈²⁻ + 2Fe²⁺ → 2SO₄²⁻ + 2Fe³⁺

2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂

Each step is between oppositely charged ions, so the activation energy is lower, and the Fe²⁺ is regenerated. Mn²⁺ behaves as an autocatalyst in the MnO₄⁻/C₂O₄²⁻ titration: it is a product of the reaction, so the reaction starts slowly and then speeds up as Mn²⁺ builds up:

4Mn²⁺ + MnO₄⁻ + 8H⁺ → 5Mn³⁺ + 4H₂O

2Mn³⁺ + C₂O₄²⁻ → 2Mn²⁺ + 2CO₂

Two habits for every question

Every transition-metal question rewards the same two habits: state the five facts for any complex, and run the colour chain for any colour change.

🔗 Next, put this to work identifying ions: the reactions of metal ions in aqueous solution (3.2.6): their acidity and the coloured precipitates they form with bases. These are the observations behind Required practical 11.