Transition metals and their chemistry

Copper sulfate is blue, potassium manganate(VII) purple, and iron catalyses the Haber process. These properties all come from one feature of the transition metals: an incomplete set of d electrons. This topic explains complex ions, their shapes and colours, the many oxidation states of vanadium and chromium, the reactions of metal aqua ions, and how transition metals act as catalysts.

What makes a transition metal

A transition metal is a d-block element that forms one or more stable ions with an incomplete d sub-shell. Across the period from scandium to zinc, electrons are added to the 3d sub-shell, but when the atoms form ions they lose their 4s electrons first.

• Chromium and copper have only one 4s electron: a half-filled (3d⁵) or full (3d¹⁰) d sub-shell is especially stable.
• Scandium is not a transition metal: its only ion, Sc³⁺, is [Ar] with an empty 3d sub-shell.
• Zinc is not a transition metal: its only ion, Zn²⁺, is [Ar] 3d¹⁰, a full d sub-shell.
• Because the 4s and 3d electrons are close in energy, successive ionisation energies rise gradually until all of them are removed, so several oxidation states are reachable.
• Characteristic properties: they form complex ions, form coloured compounds, show variable oxidation states, and act as catalysts.

Complex ions and ligands

• A complex ion is a central metal ion surrounded by ligands bonded to it by dative (coordinate) bonds.
• A ligand is a molecule or ion with a lone pair of electrons that it donates to the metal ion to form a dative covalent bond, in which both shared electrons come from the same atom.
• Coordination number: the number of dative bonds to the central metal ion (not the number of ligands).
• Monodentate ligands form one dative bond: H₂O, NH₃, Cl⁻, OH⁻, CN⁻.
• Bidentate ligands have two atoms that each donate a lone pair: ethane-1,2-diamine, H₂NCH₂CH₂NH₂ ("en"), and the ethanedioate ion, C₂O₄²⁻. Three of them give a coordination number of 6, as in [Cr(en)₃]³⁺ or [Cr(C₂O₄)₃]³⁻.
• Multidentate ligands form several bonds: EDTA⁴⁻ donates from two nitrogen atoms and four oxygen atoms, wrapping around a metal ion to form six bonds in a 1 : 1 complex. Haem is a multidentate ligand holding iron(II) in haemoglobin.
• Charge of a complex = charge of the metal ion + charges of the ligands: Cu²⁺ with four Cl⁻ gives [CuCl₄]²⁻.

Shapes of complex ions

Cisplatin and haemoglobin

Cisplatin is a square planar platinum(II) complex with the two ammonia ligands next to each other (cis) and the two chloride ligands next to each other. Its isomer with them opposite each other, transplatin, has no anticancer action. In the body a chloride ligand is replaced by water, [Pt(NH₃)₂Cl₂] + H₂O → [Pt(NH₃)₂Cl(H₂O)]⁺ + Cl⁻, and the platinum then forms dative bonds to nitrogen atoms on neighbouring bases in DNA, kinking the strand so the cancer cell cannot copy its DNA and divide. It also binds to the DNA of healthy cells, which causes side effects such as hair loss.

Haemoglobin carries oxygen as a ligand bonded to Fe²⁺ in haem. Carbon monoxide is toxic because it forms a much stronger dative bond to the iron, replacing oxygen in a ligand exchange that is hard to reverse, so less oxygen can be carried.

Why transition metal compounds are coloured

In an isolated ion the five 3d orbitals have the same energy. When ligands bond to the ion, their lone pairs repel the d electrons, and in an octahedral complex the d orbitals split into two groups at different energies, separated by ΔE. A d electron can absorb a photon of visible light whose energy exactly matches ΔE and jump to the upper group:

\[ \Delta E = hf \]

The light that is absorbed is removed from white light, and the colour we see is the complementary colour of the light that is transmitted or reflected. Hydrated copper(II) ions absorb in the orange-red region, so they look blue.

• What changes the colour: anything that changes ΔE: a different ligand, a different coordination number (and so shape), or a different oxidation state of the metal.
• Colourless ions: Sc³⁺ (3d⁰) has no d electrons to promote, and Zn²⁺ and Cu⁺ (3d¹⁰) have no space in the upper d orbitals, so they absorb no visible light.

Measuring concentration by colorimetry

The amount of light a coloured solution absorbs is proportional to the concentration of the coloured species. A colorimeter (spectrophotometer) uses a filter of the complementary colour, the colour the solution absorbs most strongly. If a complex is only pale, first add a ligand that forms a more intensely coloured complex. Make up solutions of known concentration, measure their absorbance, and plot a calibration curve; then measure the unknown and read its concentration from the graph.

Variable oxidation states

The 4s and 3d electrons are similar in energy, so transition metals can lose different numbers of them. Some general trends: high oxidation states such as manganese(VII) in MnO₄⁻ are oxidising agents; low ones such as V²⁺ and Fe²⁺ are reducing agents; and the +2 state becomes more stable relative to +3 across the period.

Vanadium

Adding zinc to acidified vanadate(V) reduces it step by step, through yellow, then green (a mixture of yellow and blue), blue, green and finally violet. Electrode potentials explain how far a reducing agent can go:

• VO₂⁺ + 2H⁺ + e⁻ ⇌ VO²⁺ + H₂O, +1.00 V; VO²⁺ + 2H⁺ + e⁻ ⇌ V³⁺ + H₂O, +0.34 V; V³⁺ + e⁻ ⇌ V²⁺, −0.26 V.
• Zinc (Zn²⁺/Zn −0.76 V) is more negative than all three, so it reduces vanadium all the way to V²⁺.
• Tin (Sn²⁺/Sn −0.14 V) is more negative than the first two but not the third, so it stops at V³⁺.

Chromium

• Reducing dichromate(VI): zinc and hydrochloric acid reduce orange Cr₂O₇²⁻ to green Cr³⁺, then to blue Cr²⁺: Cr₂O₇²⁻ + 14H⁺ + 3Zn → 2Cr³⁺ + 7H₂O + 3Zn²⁺, then 2Cr³⁺ + Zn → 2Cr²⁺ + Zn²⁺. Air must be kept out (for example with a stream of hydrogen), because oxygen oxidises Cr²⁺ straight back to Cr³⁺.
• Iron(II) is a weaker reducing agent: E^⦵(Fe³⁺/Fe²⁺) = +0.77 V lies between E^⦵(Cr₂O₇²⁻/Cr³⁺) = +1.33 V and E^⦵(Cr³⁺/Cr²⁺) = −0.41 V, so Fe²⁺ reduces dichromate only as far as Cr³⁺. Cr₂O₇²⁻ + 14H⁺ + 6Fe²⁺ → 2Cr³⁺ + 7H₂O + 6Fe³⁺ is used in titrations (with an indicator, as the colour change is not sharp enough on its own).
• Oxidising chromium(III) is easier in alkaline solution, because removing electrons from a negative ion is easier. With excess NaOH, Cr³⁺ forms green [Cr(OH)₆]³⁻, which hydrogen peroxide oxidises to yellow chromate(VI): 2[Cr(OH)₆]³⁻ + 3H₂O₂ → 2CrO₄²⁻ + 2OH⁻ + 8H₂O.
• Chromate and dichromate are linked by an acid–base equilibrium, not a redox one (chromium is +6 in both): 2CrO₄²⁻ + 2H⁺ ⇌ Cr₂O₇²⁻ + H₂O. Adding acid turns yellow chromate orange; adding alkali removes H⁺ and turns it back.

Reactions of metal aqua ions

A Lewis acid accepts an electron pair; a Lewis base donates one. In a complex the ligands are Lewis bases and the metal ion is a Lewis acid. Dissolved transition metal salts form hexaaqua ions, [M(H₂O)₆]²⁺ or [M(H₂O)₆]³⁺.

Acidity of aqua ions

The metal ion pulls electron density from the oxygen of each water ligand, weakening its O–H bonds, so a proton can be lost to the solvent:

[Fe(H₂O)₆]³⁺ + H₂O ⇌ [Fe(H₂O)₅(OH)]²⁺ + H₃O⁺

3+ ions are smaller and more highly charged than 2+ ions, so they have a greater charge density and polarise the water ligands more strongly. Solutions of M³⁺ ions are therefore noticeably acidic (around pH 2 to 3), while M²⁺ solutions are only very weakly acidic. This hydrolysis is why iron(III) solutions look yellow-brown rather than pale violet.

Deprotonation: hydroxide precipitates

Hydroxide ions and ammonia act as Brønsted–Lowry bases, removing protons from water ligands one at a time until a neutral complex forms, which precipitates:

• [Cu(H₂O)₆]²⁺ + 2OH⁻ → [Cu(OH)₂(H₂O)₄] + 2H₂O (blue precipitate)
• [Cu(H₂O)₆]²⁺ + 2NH₃ → [Cu(OH)₂(H₂O)₄] + 2NH₄⁺ (the same precipitate with a little ammonia)
• [Fe(H₂O)₆]³⁺ + 3OH⁻ → [Fe(OH)₃(H₂O)₃] + 3H₂O (orange-brown precipitate)

Amphoteric hydroxides

The hydroxides of chromium(III) and zinc dissolve in excess sodium hydroxide, because further deprotonation gives a soluble negative ion; they also dissolve in acid. A hydroxide that reacts with both acids and alkalis is amphoteric:

• [Cr(OH)₃(H₂O)₃] + 3OH⁻ → [Cr(OH)₆]³⁻ + 3H₂O (dark green solution)
• [Zn(OH)₂(H₂O)₄] + 2OH⁻ → [Zn(OH)₄]²⁻ + 4H₂O (colourless solution)
• With acid: [Cr(OH)₃(H₂O)₃] + 3H⁺ → [Cr(H₂O)₆]³⁺.

Ligand exchange

• With excess ammonia, NH₃ replaces water (and hydroxide) ligands on copper, cobalt, nickel, chromium(III) and zinc, dissolving the precipitate. NH₃ and H₂O are similar in size and both uncharged, so the coordination number stays 6 (exchange may be incomplete, as with copper): [Cu(OH)₂(H₂O)₄] + 4NH₃ → [Cu(NH₃)₄(H₂O)₂]²⁺ + 2H₂O + 2OH⁻ (deep blue).
• With concentrated HCl (a high concentration of Cl⁻), chloride replaces water. Cl⁻ is larger than H₂O, so only four fit and the coordination number falls from 6 to 4, giving tetrahedral complexes: [Cu(H₂O)₆]²⁺ + 4Cl⁻ ⇌ [CuCl₄]²⁻ + 6H₂O (yellow-green); [Co(H₂O)₆]²⁺ + 4Cl⁻ ⇌ [CoCl₄]²⁻ + 6H₂O (pink to blue). Adding water reverses these.

Stability of complexes and the chelate effect

Replacing monodentate ligands with bidentate or multidentate ones gives a much more stable complex. This chelate effect is explained by entropy:

• [Cu(H₂O)₆]²⁺ + EDTA⁴⁻ → [Cu(EDTA)]²⁻ + 6H₂O: 2 particles become 7.
• [Co(NH₃)₆]²⁺ + 3en → [Co(en)₃]²⁺ + 6NH₃: 4 particles become 7.
• More particles means more ways to arrange the energy, so ΔS_system is positive. The same number and type of dative bonds are broken and made (M–O or M–N), so ΔH is close to zero and ΔS_surroundings is small. ΔS_total is therefore positive and the chelate complex is favoured.
• Uses of EDTA: it binds metal ions so tightly that it is used to remove toxic heavy-metal ions from rivers and from the body, and to soften water in shampoos by trapping Ca²⁺.

Working out formulae with silver nitrate

Chloride can be a ligand bonded to the metal or a free ion outside the complex. Silver nitrate precipitates only the free chloride ions as AgCl, so the mole ratio shows how many are outside. The three isomers of formula CrCl₃·6H₂O behave differently:

Transition metals as catalysts

A catalyst speeds up a reaction by providing an alternative route with a lower activation energy, and is unchanged at the end. Transition metals are good catalysts because they can change oxidation state easily, and because they can use their 3d and 4s electrons to form weak bonds to reactants on their surface.

Heterogeneous catalysis

A heterogeneous catalyst is in a different phase from the reactants, usually a solid with gases or liquids. The reaction happens on its surface:

1. Reactant molecules are adsorbed onto active sites on the surface, forming weak bonds to the metal.
2. Bonds within the reactants are weakened, and the molecules are held close together in a favourable orientation, so they react more easily.
3. The products are desorbed from the surface, freeing the sites.

• Strength of adsorption must be just right: tungsten adsorbs too strongly, so products are not released; silver too weakly, so too few reactant molecules stick. Nickel and platinum are in between and make good catalysts.
• Surface area: the catalyst is spread thinly on a support, such as a ceramic honeycomb, to give a large surface area from a small amount of expensive metal.
• Advantages: solid catalysts are easy to separate from the products and suit continuous processes.
• Examples: iron in the Haber process; vanadium(V) oxide in the Contact process, which changes oxidation state and changes back: SO₂ + V₂O₅ → SO₃ + V₂O₄, then 2V₂O₄ + O₂ → 2V₂O₅; platinum, palladium and rhodium in catalytic converters, which remove carbon monoxide, nitrogen monoxide and unburnt hydrocarbons, for example 2CO + 2NO → 2CO₂ + N₂.
• Developing better catalysts: ethanoic acid is now made from methanol and carbon monoxide, CH₃OH + CO → CH₃COOH, with 100% atom economy, using an iridium–iodide catalyst. Successive catalysts have allowed milder conditions and better selectivity; such advances are published in peer-reviewed journals so other scientists can check them.

Homogeneous catalysis

A homogeneous catalyst is in the same phase as the reactants. It works by forming an intermediate, usually by changing oxidation state and then changing back.

Autocatalysis

In autocatalysis a product catalyses the reaction. In the titration of ethanedioate with manganate(VII), 2MnO₄⁻ + 16H⁺ + 5C₂O₄²⁻ → 2Mn²⁺ + 10CO₂ + 8H₂O, the reaction starts slowly (again two negative ions must collide), then speeds up as Mn²⁺ builds up, because Mn²⁺ provides a faster route:

• 4Mn²⁺ + MnO₄⁻ + 8H⁺ → 5Mn³⁺ + 4H₂O
• 2Mn³⁺ + C₂O₄²⁻ → 2Mn²⁺ + 2CO₂
• The rate can be followed with a colorimeter measuring the purple colour, which does not disturb the reaction, or by titrating samples.

Quick checks

Interactive diagrams