28.2 Chemical properties of first-row transition elements

Ligands can replace one another around a metal ion. Water, ammonia, hydroxide, chloride and other ligands donate lone pairs and alter charge, colour and geometry.

Balance the metal, ligands and charge in the equation. Excess ligand or precipitating hydroxide can drive a substitution and may produce a different coordination number.

Cu²⁺(aq) + 4NH₃(aq) ⇌ [Cu(NH₃)₄]²⁺ changes the ligand environment and colour; adding OH⁻ can instead form Cu(OH)₂ precipitate.

A colour change does not by itself prove oxidation-state change; ligand substitution can change colour while Cu remains +2.

A ligand is an ion or molecule with at least one lone pair that forms a dative covalent bond to a metal atom or ion. The donor atom is the atom directly attached to the metal.

Water and ammonia donate oxygen or nitrogen lone pairs; chloride and cyanide donate anionic lone pairs. The ligand’s charge contributes to the complex charge but does not change the donor definition.

In [Co(NH₃)₆]³⁺, each NH₃ donates one nitrogen lone pair and the six ligands are neutral.

A ligand is not any surrounding ion, and a coordinate bond is not formed by the metal donating both electrons.

A monodentate ligand binds through one donor atom, a bidentate ligand through two, and a polydentate ligand through several donor atoms.

H₂O, NH₃, Cl⁻ and CN⁻ are monodentate; en and ethanedioate are bidentate; EDTA⁴⁻ is polydentate. Denticity affects coordination number and chelate stability.

One en ligand occupies two coordination sites, so three en ligands can surround a metal in an octahedral complex.

Denticity is not the same as the ligand’s charge or the number of ligands written in the formula.

A complex is a molecule or ion in which a central metal atom or ion is bonded to surrounding ligands through coordinate bonds. The bracketed species has its own overall charge.

Separate the complex ion from counter-ions outside the brackets. Count ligand donor atoms to determine coordination number, then check total charge.

In [CoCl₄]²⁻, Co is central, four chloride ligands coordinate, and the bracketed species carries −2 charge.

The counter-ions are not ligands inside the complex, and coordination number is not always equal to the number of written ligand molecules for polydentate ligands.

Common complex geometries are linear (2 sites, 180°), square planar (4, 90°), tetrahedral (4, about 109.5°) and octahedral (6, 90° between adjacent sites).

Geometry depends on coordination number, ligand size, metal and electronic configuration. Use the stated complex rather than assigning a shape from charge alone.

[Ag(NH₃)₂]⁺ is linear; [PtCl₄]²⁻ square planar; [CoCl₄]²⁻ tetrahedral; [Cu(H₂O)₆]²⁺ octahedral (often distorted).

Four-coordinate does not uniquely mean tetrahedral—square-planar complexes are also possible.

Coordination number is the number of donor atoms directly bonded to the central metal, not always the number of ligand molecules. The complex charge is metal charge plus all ligand charges.

Choose ligands to fill the stated coordination number, then add charges algebraically. Denticity matters: one bidentate ligand occupies two sites.

A +2 metal with four neutral NH₃ ligands gives [M(NH₃)₄]²⁺; a +3 metal with two oxalate ions (2− each) gives [M(C₂O₄)₂]⁻.

Do not use ligand count as coordination number for en or EDTA, and do not confuse oxidation state with overall complex charge.

Ligand exchange is a substitution at a metal centre: one ligand leaves as another donates a lone pair. The metal oxidation state may stay the same while colour and geometry change.

Use the ligand and stoichiometry to balance the equation. Copper(II) and cobalt(II) complexes show exchange with water, ammonia, hydroxide or chloride.

Adding excess NH₃ to [Cu(H₂O)₆]²⁺ forms a deep-blue ammine complex; adding OH⁻ can produce Cu(OH)₂ instead. The observations depend on ligand amount and medium.

Ligand exchange is not automatically redox and a colour change does not prove the metal changed oxidation state.

Write the relevant reduction half-equations, select the more positive reduction as the cathode, and calculate E°cell = E°cathode − E°anode. A positive value supports the written standard direction.

Transition-metal ions can act as oxidising or reducing agents depending on their pair and oxidation state. Balance electrons after choosing the direction.

If MnO₄⁻/Mn²⁺ is more positive than Fe³⁺/Fe²⁺ in acid, permanganate can oxidise Fe²⁺ to Fe³⁺ while being reduced.

Do not rank ions without specifying the half-equations or assume a reaction is feasible under non-standard concentrations from E° alone.

In acid, MnO₄⁻ is reduced to Mn²⁺; oxalate, Fe²⁺ or I⁻ are oxidised according to their half-equations. Combine half-equations so electrons, atoms and charge balance.

Use stoichiometric ratios to calculate titres or amounts. For Cu²⁺ + I⁻, iodine forms and can be titrated with thiosulfate in the full analytical method.

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O; Fe²⁺ → Fe³⁺ + e⁻, so one permanganate reacts with five Fe²⁺.

Do not balance oxygen with OH⁻ in an acidic equation without converting correctly, and do not ignore the reaction medium.

For any supplied redox pair, identify oxidation and reduction, balance each half-equation, cancel electrons and use the resulting mole ratio for calculations.

Use E° or the stated conditions to choose direction, then verify mass and charge conservation. Unfamiliar species do not require a new method.

If X²⁺ + 2e⁻ → X and Y → Y³⁺ + 3e⁻, the least common electron number is six, giving 3X²⁺ + 2Y → 3X + 2Y³⁺.

Do not force a familiar coefficient pattern onto a new oxidation state or calculate mass before finding the electron ratio.