OxfordAQA A-Level Chemistry: Transition metals (9620)

This study guide teaches section 3.2.5, Transition metals, from OxfordAQA’s International AS and A-level Chemistry specification (9620), Version 5.3, which applies to exams from May/June 2020. It covers 3.2.5.1 to 3.2.5.6. This is International A2 content, so everything here is International A-level only. Unit 3 assesses it, and the synoptic Unit 5 can draw on any content.

Pair it with the revision notes and practice questions. Related pages: electrode potentials, redox equations and thermodynamics. For the wider course: hub page, printable checklist, free diagnostics.

What this section covers

All rows are International A-level only; no required practical is listed here.

Spec ref	You need to be able to
3.2.5.1	Link properties to an incomplete sub-level; define ligand, complex, co-ordination number
3.2.5.2	Ligand exchange; bidentate and multidentate ligands; haem; the chelate effect
3.2.5.3	Four shapes; cis–trans and optical isomerism; cisplatin; Tollens’ reagent
3.2.5.4	Explain colour; ΔE = hν = hc/λ; colorimetry
3.2.5.5	Vanadium; pH and ligand effects; Tollens’ test; MnO₄⁻ titration calculations
3.2.5.6	Heterogeneous and homogeneous catalysis, with equations

3.2.5.1 General properties of transition metals

From titanium to copper, each element has an incomplete d sub-level in its atom or in at least one of its ions. Four properties follow from it:

• they form complexes
• they form coloured ions
• they have variable oxidation states
• they show catalytic activity.

Three definitions underpin the whole section:

• Ligand: an ion or molecule that bonds to a transition metal ion by donating an electron pair, making a co-ordinate bond.
• Complex: ligands arranged around a central metal atom or ion.
• Co-ordination number: how many co-ordinate bonds the central metal atom or ion forms.

Worked example 1: reading a complex

Take [Mn(H₂O)₆]²⁺.

1. Ligand: H₂O, donating a lone pair from oxygen.
2. Co-ordination number: six water ligands, one bond each, so 6.
3. Oxidation state: water is neutral, so manganese is +2.
4. Mn²⁺ is [Ar] 3d⁵ (4s electrons go first): an incomplete sub-level.

3.2.5.2 Substitution reactions

Three monodentate ligands appear here, each forming one co-ordinate bond: H₂O, NH₃ and Cl⁻.

Swapping water and ammonia

NH₃ and H₂O are uncharged and close in size, so one replaces the other and the co-ordination number stays at 6:

6NH₃(aq) + [Co(H₂O)₆]²⁺(aq) → [Co(NH₃)₆]²⁺(aq) + 6H₂O(l)

Substitution can be incomplete. Copper(II) swaps only four waters and stays six-co-ordinate:

4NH₃(aq) + [Cu(H₂O)₆]²⁺(aq) → [Cu(NH₃)₄(H₂O)₂]²⁺(aq) + 4H₂O(l)

Swapping water for chloride

Cl⁻ is bigger than NH₃ or H₂O, so fewer fit round the metal and the co-ordination number can fall from 6 to 4. With concentrated hydrochloric acid:

4Cl⁻(aq) + [Co(H₂O)₆]²⁺(aq) → [CoCl₄]²⁻(aq) + 6H₂O(l)
4Cl⁻(aq) + [Cu(H₂O)₆]²⁺(aq) → [CuCl₄]²⁻(aq) + 6H₂O(l)
4Cl⁻(aq) + [Fe(H₂O)₆]³⁺(aq) → [FeCl₄]⁻(aq) + 6H₂O(l)

Ligands with more than one donor atom

• Bidentate ligands bond through two atoms. The specification’s examples are ethane-1,2-diamine, H₂NCH₂CH₂NH₂ (each nitrogen donates), and ethanedioate, C₂O₄²⁻ (two oxygens donate).
• Multidentate ligands bond through several atoms. EDTA⁴⁻ uses six.

Haem is an iron(II) complex in which the ligand is multidentate. In haemoglobin, O₂ forms a co-ordinate bond to Fe(II), so blood can carry oxygen. Carbon monoxide is toxic: it bonds to Fe(II) in place of that O₂.

The chelate effect

Bidentate and multidentate ligands displace monodentate ones from a complex. This is the chelate effect. Explain it through the balance of enthalpy and entropy:

• The donor atoms barely change and the bond count stays the same, so ΔH is close to zero.
• One chelating ligand frees several small ligands, so the number of particles rises and ΔS is positive.
• With ΔH small and ΔS positive, ΔG = ΔH − TΔS is negative, so the substitution is feasible.

Worked example 2: two diamine ligands on copper

[Cu(H₂O)₆]²⁺(aq) + 2H₂NCH₂CH₂NH₂(aq) → [Cu(H₂NCH₂CH₂NH₂)₂(H₂O)₂]²⁺(aq) + 4H₂O(l)

1. Particles: 1 + 2 = 3 on the left, 1 + 4 = 5 on the right. Entropy increases.
2. Suppose a data book gave ΔH = +3.5 kJ mol⁻¹ and ΔS = +48 J K⁻¹ mol⁻¹ for this change (values chosen for practice).
3. Convert ΔS: +0.048 kJ K⁻¹ mol⁻¹.
4. At 298 K: ΔG = 3.5 − (298 × 0.048) = −10.8 kJ mol⁻¹.

ΔG is negative although ΔH is positive: entropy drives it.

3.2.5.3 Shapes of complex ions

Ligands	Shape	Bond angle	Example
Small, such as H₂O and NH₃	octahedral	90°	[Co(NH₃)₆]²⁺
Larger, such as Cl⁻	tetrahedral	109.5°	[CoCl₄]²⁻
Four ligands in one plane	square planar	90°	cisplatin
Ag⁺ with two NH₃	linear	180°	[Ag(NH₃)₂]⁺, Tollens’ reagent

Isomerism in complexes

• Octahedral with monodentate ligands: cis–trans isomerism, which is a special case of E–Z isomerism. Two identical ligands sit either at 90° to each other (cis) or opposite each other at 180° (trans).
• Octahedral with bidentate ligands: optical isomerism; three bidentate ligands give non-superimposable mirror images.
• Square planar: cis–trans isomerism again. Cisplatin is the cis isomer of [Pt(NH₃)₂Cl₂]: chloride ligands side by side.

Worked example 3: building a complex from a description

Cobalt(III) bonds to two ethane-1,2-diamine ligands and two chloride ions.

1. Formula: [Co(H₂NCH₂CH₂NH₂)₂Cl₂], charge still to find.
2. Co-ordination number: 2 × 2 + 2 × 1 = 6, so octahedral.
3. Charge: +3 + 2(0) + 2(−1) = +1.
4. Isomers: the two Cl⁻ ligands can be at 90° (cis) or 180° (trans). The cis form also has no plane of symmetry, so it exists as a pair of optical isomers.

3.2.5.4 Formation of coloured ions

Colour is one way to identify a transition metal ion. The ion absorbs some visible wavelengths; the rest are transmitted or reflected, and that is what you see.

Absorbing a photon promotes a d electron: it leaves its ground state for an excited state. The energy gap matches the photon:

ΔE = hν = hc/λ

A change in oxidation state, co-ordination number or ligand alters ΔE. A different wavelength is absorbed, so the colour changes. Pink [Co(H₂O)₆]²⁺ turning blue as it becomes [CoCl₄]²⁻ involves both a new ligand and a new co-ordination number.

Worked example 4: from frequency to wavelength

A complex absorbs light of frequency 5.20 × 10¹⁴ Hz. Take h = 6.63 × 10⁻³⁴ J s, L = 6.02 × 10²³ mol⁻¹ and c = 3.00 × 10⁸ m s⁻¹.

1. ΔE = hν = 6.63 × 10⁻³⁴ × 5.20 × 10¹⁴ = 3.45 × 10⁻¹⁹ J for one ion.
2. Per mole: 3.4476 × 10⁻¹⁹ × 6.02 × 10²³ = 2.08 × 10⁵ J, so 208 kJ mol⁻¹.
3. λ = c/ν = 3.00 × 10⁸ ÷ 5.20 × 10¹⁴ = 5.77 × 10⁻⁷ m = 577 nm.

Spectroscopy and the colorimeter

Spectroscopy uses this absorption; a simple colorimeter finds the concentration of a coloured ion:

1. Choose a filter giving the colour of light the ion absorbs most.
2. Zero the instrument with solvent only.
3. Plot absorbance against concentration for several standards.
4. Read the unknown’s concentration off this graph.

Worked example 5: nickel(II) by calibration

Standards of Ni²⁺(aq) at 0.050, 0.100, 0.150 and 0.200 mol dm⁻³ read 0.11, 0.22, 0.33 and 0.44. A 10.0 cm³ sample was diluted to 50.0 cm³ and read 0.29.

1. The line passes through the origin with gradient 0.11 ÷ 0.050 = 2.2 dm³ mol⁻¹.
2. Diluted solution: 0.29 ÷ 2.2 = 0.132 mol dm⁻³.
3. Dilution factor 50.0 ÷ 10.0 = 5, so the sample was 0.659 mol dm⁻³.

3.2.5.5 Variable oxidation states

Vanadium is the specification’s example. Zinc in acid reduces vanadate(V) one step at a time:

Oxidation state of V	Species	Colour
+5	VO₂⁺	yellow
+4	VO²⁺	blue
+3	V³⁺	green
+2	V²⁺	violet

VO₂⁺(aq) + 2H⁺(aq) + e⁻ → VO²⁺(aq) + H₂O(l)
VO²⁺(aq) + 2H⁺(aq) + e⁻ → V³⁺(aq) + H₂O(l)
V³⁺(aq) + e⁻ → V²⁺(aq)
Zn(s) → Zn²⁺(aq) + 2e⁻

A green tint also shows up between yellow and blue, where yellow VO₂⁺ and blue VO²⁺ are both present.

pH and ligand change the redox potential

How easily an ion drops to a lower oxidation state depends on pH and on the ligand.

• pH: MnO₄⁻ + 8H⁺ + 5e⁻ ⇌ Mn²⁺ + 4H₂O uses H⁺, so manganate(VII) oxidises more strongly at low pH.
• Ligand: Ag⁺/Ag has E⦵ = +0.80 V, but [Ag(NH₃)₂]⁺/Ag is only +0.37 V. Ammonia makes silver(I) harder to reduce.

This mild oxidising agent is Tollens’ reagent. Aldehydes reduce it to silver; ketones do not:

[Ag(NH₃)₂]⁺(aq) + e⁻ → Ag(s) + 2NH₃(aq)

Titrations with manganate(VII)

MnO₄⁻(aq) + 8H⁺(aq) + 5Fe²⁺(aq) → Mn²⁺(aq) + 5Fe³⁺(aq) + 4H₂O(l)                (1 : 5)
2MnO₄⁻(aq) + 16H⁺(aq) + 5C₂O₄²⁻(aq) → 2Mn²⁺(aq) + 10CO₂(g) + 8H₂O(l)          (2 : 5)

The titrant is purple and is decolourised while the analyte remains, so the first lasting pink colour marks the end point. Acidify with dilute sulfuric acid; chloride would be oxidised by MnO₄⁻ and give a titre that is too large. Ethanedioate titrations are run warm because they start slowly.

Worked example 6: iron in a lawn tonic

2.40 g of an iron(II) lawn tonic goes into dilute sulfuric acid, and the solution is made up to 250 cm³. Titrating 25.0 cm³ of it takes 9.40 cm³ of 0.0185 mol dm⁻³ KMnO₄. Ar(Fe) = 55.8.

1. n(MnO₄⁻) = 0.0185 × 9.40 ÷ 1000 = 1.739 × 10⁻⁴ mol.
2. n(Fe²⁺) = 5 × 1.739 × 10⁻⁴ = 8.695 × 10⁻⁴ mol in 25.0 cm³, so 8.695 × 10⁻³ mol in 250 cm³.
3. Mass of iron = 8.695 × 10⁻³ × 55.8 = 0.485 g.
4. Percentage = 0.485 ÷ 2.40 × 100 = 20.2%.

3.2.5.6 Catalysts

Heterogeneous catalysts are in a different phase from the reactants; reaction happens at active sites on the surface. A support medium maximises surface area and minimises cost. Impurities can poison active sites, cutting efficiency, which costs money.

• V₂O₅ is the heterogeneous catalyst in the Contact process.
• Fe is the heterogeneous catalyst in the Haber process.

A homogeneous catalyst shares the reactants’ phase, and the reaction goes through an intermediate species.

Variable oxidation states let the catalyst take electrons in one step and return them in the next, giving a lower-activation-energy route.

V₂O₅ in the Contact process

V₂O₅(s) + SO₂(g) → V₂O₄(s) + SO₃(g)          V goes from +5 to +4
V₂O₄(s) + ½O₂(g) → V₂O₅(s)                    V goes from +4 back to +5

Fe²⁺ with peroxodisulfate and iodide

Both reactants are negative ions, which repel. Fe²⁺ gives two steps between oppositely charged ions:

2Fe²⁺(aq) + S₂O₈²⁻(aq) → 2Fe³⁺(aq) + 2SO₄²⁻(aq)
2Fe³⁺(aq) + 2I⁻(aq) → 2Fe²⁺(aq) + I₂(aq)

Mn²⁺ as an autocatalyst

Mn²⁺ is a product that catalyses its own formation. The rate starts low, rises as Mn²⁺ builds up, then falls as reactants run out.

MnO₄⁻(aq) + 4Mn²⁺(aq) + 8H⁺(aq) → 5Mn³⁺(aq) + 4H₂O(l)
2Mn³⁺(aq) + C₂O₄²⁻(aq) → 2Mn²⁺(aq) + 2CO₂(g)

Common errors

• Giving the number of ligands as the co-ordination number of a chelate.
• Explaining the chelate effect with “stronger bonds” rather than more particles.
• Using 1 : 5 for MnO₄⁻ : C₂O₄²⁻ (it is 2 : 5).
• Confusing Fe metal (heterogeneous, Haber) with Fe²⁺(aq) (homogeneous).

Official syllabus

Source: OxfordAQA’s International AS and A-level Chemistry specification (9620), Version 5.3, for exams from May/June 2020. Section 3.2.5 Transition metals (International A2), subsections 3.2.5.1 to 3.2.5.6.