General characteristic chemical properties of the first set of transition elements, titanium to copper

· Transition elements form complexes by reacting with ligands through coordinate / dative covalent bonds.
· A ligand = a species with a lone pair of electrons that forms a dative covalent bond to a central metal atom or ion.
· A complex = a molecule or ion containing a central metal atom/ion surrounded by ligands.
· Transition metal ions often have vacant, energetically accessible d orbitals, allowing ligand donation and complex formation.
· In equations, complexes are written in square brackets, with the overall charge outside: e.g. [Cu(H₂O)₆]²⁺.

This diagram links coordination number to the common shapes of transition metal complexes. It helps distinguish octahedral, tetrahedral and square planar arrangements, which are central to CIE 28.2. Source

Ligands: monodentate, bidentate and polydentate

· Monodentate ligand = donates one lone pair to form one coordinate bond.
· Required examples: H₂O, NH₃, Cl⁻, CN⁻.
· Bidentate ligand = donates two lone pairs to form two coordinate bonds to the same metal ion.
· Required examples: 1,2-diaminoethane / en / H₂NCH₂CH₂NH₂ and ethanedioate ion / C₂O₄²⁻.
· Polydentate ligand = donates more than two lone pairs to form several coordinate bonds.
· Required example: EDTA⁴⁻.
· Exam trap: number of ligands ≠ always coordination number; bidentate and polydentate ligands form multiple coordinate bonds per ligand.

This clean diagram shows how a tetrahedral ligand arrangement affects the metal d orbitals. Although orbital splitting belongs more directly to colour, the tetrahedral arrangement is important for recognising four-coordinate complexes in 28.2. Source

Coordination number, formula and charge

· Coordination number = the number of coordinate bonds formed between ligands and the central metal atom/ion.
· For monodentate ligands, coordination number = number of ligands.
· For bidentate ligands, each ligand contributes 2 coordinate bonds.
· For EDTA⁴⁻, one ligand can form six coordinate bonds.
· Overall complex charge = metal ion charge + total ligand charges.
· Neutral ligands: H₂O = 0, NH₃ = 0, en = 0.
· Negative ligands: Cl⁻ = –1, CN⁻ = –1, C₂O₄²⁻ = –2, EDTA⁴⁻ = –4.
· Example: Cu²⁺ + 4Cl⁻ → [CuCl₄]²⁻ because +2 + 4(–1) = –2.
· Example: Co²⁺ + 6NH₃ → [Co(NH₃)₆]²⁺ because NH₃ is neutral, so charge remains +2.

Shapes and bond angles of complexes

· Linear: coordination number 2, bond angle 180°.
· Example: [Ag(NH₃)₂]⁺.
· Tetrahedral: coordination number 4, bond angle 109.5°.
· Example: [CuCl₄]²⁻ or [CoCl₄]²⁻.
· Square planar: coordination number 4, bond angle 90°.
· Example: [Pt(NH₃)₂Cl₂].
· Octahedral: coordination number 6, bond angle 90°.
· Example: [Cu(H₂O)₆]²⁺, [Co(H₂O)₆]²⁺, [Co(NH₃)₆]²⁺.
· Exam trap: coordination number 4 may be tetrahedral or square planar, so use the information given in the question.

This diagram is a useful A-Level visual summary of linear, tetrahedral, square planar and octahedral complex shapes. It directly supports exam questions asking for shape, bond angle and coordination number. Source

Ligand exchange reactions

· Ligand exchange = one ligand in a complex is replaced by another ligand.
· Occurs because ligands form coordinate bonds to the central metal ion.
· Ligand exchange may change the formula, coordination number, geometry and sometimes the colour of the complex.
· H₂O and NH₃ are neutral ligands, so exchanging between them usually keeps the overall charge the same.
· Cl⁻ and OH⁻ are negative ligands, so they can change the overall charge or cause precipitation.
· In ligand exchange equations, balance atoms and charge carefully.

Copper(II) complexes with water, ammonia, hydroxide and chloride

· In water, copper(II) commonly exists as [Cu(H₂O)₆]²⁺, an octahedral complex.
· With limited hydroxide:
[Cu(H₂O)₆]²⁺ + 2OH⁻ → Cu(H₂O)₄(OH)₂ + 2H₂O.
· This forms a copper(II) hydroxide complex precipitate, commonly written as Cu(OH)₂(s) in simpler ionic equations.
· With excess ammonia:
[Cu(H₂O)₆]²⁺ + 4NH₃ ⇌ [Cu(NH₃)₄(H₂O)₂]²⁺ + 4H₂O.
· Coordination number remains 6, but four water ligands are replaced by NH₃.
· With excess chloride ions:
[Cu(H₂O)₆]²⁺ + 4Cl⁻ ⇌ [CuCl₄]²⁻ + 6H₂O.
· Coordination number changes from 6 to 4 and geometry changes from octahedral to tetrahedral.

The image shows the visible change when NH₃ ligands replace H₂O ligands around Cu²⁺. It is a clear example of ligand exchange in copper(II) complexes. Source

Cobalt(II) complexes with water, ammonia, hydroxide and chloride

· In water, cobalt(II) commonly exists as [Co(H₂O)₆]²⁺, an octahedral complex.
· With limited hydroxide:
[Co(H₂O)₆]²⁺ + 2OH⁻ → Co(H₂O)₄(OH)₂ + 2H₂O.
· This forms a cobalt(II) hydroxide complex precipitate, often simplified as Co(OH)₂(s).
· With ammonia:
[Co(H₂O)₆]²⁺ + 6NH₃ ⇌ [Co(NH₃)₆]²⁺ + 6H₂O.
· Coordination number remains 6 and the complex remains octahedral.
· With excess chloride ions:
[Co(H₂O)₆]²⁺ + 4Cl⁻ ⇌ [CoCl₄]²⁻ + 6H₂O.
· Coordination number changes from 6 to 4 and geometry changes from octahedral to tetrahedral.

Using E⦵ values to predict feasibility

· Use standard electrode potentials, E⦵, to predict whether a redox reaction is feasible.
· More positive E⦵ = species is more easily reduced = stronger oxidising agent.
· More negative E⦵ = reduced form is more easily oxidised = stronger reducing agent.
· For a proposed redox reaction:
E⦵cell = E⦵reduction − E⦵oxidation.
· If E⦵cell is positive, the reaction is feasible under standard conditions.
· If E⦵cell is negative, the reaction is not feasible under standard conditions.
· Always reverse the half-equation for the species being oxidised, but do not change the sign until using it as the oxidation side in E⦵cell.
· Exam phrase: a positive E⦵cell means the reaction is thermodynamically feasible, but it may still be slow due to kinetic factors.

Redox calculations: key half-equations

· In acidic solution, manganate(VII) is reduced as:
MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O.
· Ethanedioate ions are oxidised as:
C₂O₄²⁻ → 2CO₂ + 2e⁻.
· Iron(II) ions are oxidised as:
Fe²⁺ → Fe³⁺ + e⁻.
· Iodide ions are oxidised as:
2I⁻ → I₂ + 2e⁻.
· Copper(II) ions are reduced by iodide to form CuI(s):
2Cu²⁺ + 4I⁻ → 2CuI(s) + I₂.
· For calculations: convert volume to dm³, use n = cV, apply the mole ratio, then calculate the required quantity.

Required redox systems and mole ratios

· MnO₄⁻ / C₂O₄²⁻ in acid solution:
2MnO₄⁻ + 5C₂O₄²⁻ + 16H⁺ → 2Mn²⁺ + 10CO₂ + 8H₂O.
· Mole ratio: 2 mol MnO₄⁻ : 5 mol C₂O₄²⁻.
· MnO₄⁻ / Fe²⁺ in acid solution:
MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O.
· Mole ratio: 1 mol MnO₄⁻ : 5 mol Fe²⁺.
· Cu²⁺ / I⁻ system:
2Cu²⁺ + 4I⁻ → 2CuI(s) + I₂.
· Mole ratio: 2 mol Cu²⁺ : 1 mol I₂.
· Acidic conditions are essential for manganate(VII) reactions to produce Mn²⁺.

Redox calculation method

· Step 1: Write the balanced redox equation or identify the mole ratio.
· Step 2: Calculate moles using n = cV, with volume in dm³.
· Step 3: Use the stoichiometric ratio to find unknown moles.
· Step 4: Convert to required quantity: concentration, mass, volume or percentage purity.
· Step 5: Quote the answer to appropriate significant figures and include units.
· For titrations, repeat titres until concordant titres are obtained, then use the mean titre.
· Exam trap: never use cm³ directly in n = cV; convert using V/dm³ = V/cm³ ÷ 1000.

Exam traps and high-grade tips

· Ligand = electron pair donor, not electron pair acceptor.
· Coordinate bond = both bonding electrons come from the ligand.
· Coordination number counts coordinate bonds, not necessarily ligands.
· For complex charge, include charges of both metal ion and ligands.
· H₂O, NH₃ and en are neutral; Cl⁻, CN⁻, C₂O₄²⁻ and EDTA⁴⁻ are charged.
· Octahedral = 6-coordinate, tetrahedral = 4-coordinate, square planar = 4-coordinate, linear = 2-coordinate.
· In manganate(VII) redox equations, remember 8H⁺ and 5e⁻ in the MnO₄⁻ half-equation.
· For E⦵cell, a positive value means feasible under standard conditions.