Transition Elements: Properties, Complexes and Redox Chemistry

This guide covers subtopics 28.1, General physical and chemical properties of the first row of transition elements, titanium to copper, and 28.2, General characteristic chemical properties of the first set of transition elements, titanium to copper, from Topic 28, Chemistry of transition elements, of Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. This is A Level content, and the first genuinely new block of inorganic chemistry beyond the AS Level Periodic Table topics.

Before studying this

This resource assumes electron configuration (including d-orbitals) from Atomic Structure: Orbitals and Ionisation Energy, oxidation numbers and half-equation balancing from Redox Processes: Oxidation Numbers and Electron Transfer, and standard electrode potentials from Electrochemistry: Electrolysis and Standard Electrode Potentials. Transition Elements: Colour, Stereoisomerism and Stability Constants continues directly from this page.

Syllabus coverage

CAMBRIDGE INTERNATIONAL AS & A LEVEL CHEMISTRY 9701 — A Level, Topic 28

28.1 General physical and chemical properties — defining a transition element as a d-block element forming one or more stable ions with incomplete d orbitals; sketching a 3d_xy and a 3d_z² orbital; understanding that transition elements have variable oxidation states, act as catalysts, form complex ions and form coloured compounds; explaining variable oxidation states in terms of the similar energy of 3d and 4s sub-shells; explaining catalytic behaviour in terms of multiple oxidation states and energetically accessible vacant d orbitals that can form dative bonds; explaining complex-ion formation in terms of energetically accessible vacant d orbitals.

28.2 General characteristic chemical properties — describing and explaining reactions of transition elements with ligands to form complexes, including copper(II) and cobalt(II) with water, ammonia, hydroxide and chloride; defining ligand, monodentate, bidentate and polydentate (with named examples); defining complex; describing complex geometry (linear, square planar, tetrahedral, octahedral); using coordination number to predict complex formula and charge; explaining ligand exchange qualitatively; predicting redox feasibility using E° values; describing and calculating using MnO₄⁻/C₂O₄²⁻, MnO₄⁻/Fe²⁺ and Cu²⁺/I⁻ reactions, and other redox systems given suitable data.

Why transition elements behave differently

A transition element is a d-block element that forms at least one stable ion with an incomplete d sub-shell. (Scandium and zinc are d-block but not transition elements by this definition — Sc³⁺ has an empty 3d sub-shell and Zn²⁺ has a full one, so neither has an incomplete d sub-shell.) The d orbitals aren’t a single spherical shape: a 3d_xy orbital has four lobes lying in the xy-plane, between the axes; a 3d_z² orbital has two lobes along the z-axis plus a “doughnut” of electron density around the middle.

Three of the four defining transition-element properties trace back to the same underlying fact: the 3d and 4s sub-shells are very close in energy.

• Variable oxidation states. Because 3d and 4s electrons are similar in energy, several can be removed at comparable energetic cost, so transition elements readily form ions of more than one oxidation state (e.g. Fe²⁺ and Fe³⁺; Cu⁺ and Cu²⁺) — unlike Group 2 metals, which are essentially always +2.
• Catalytic behaviour. Multiple accessible oxidation states let a transition-metal catalyst be oxidised and reduced within a catalytic cycle (the same logic as homogeneous catalysis by Fe²⁺/Fe³⁺), and energetically accessible vacant d orbitals let reactant molecules bind (via dative bonds) to the metal’s surface or ion, as in heterogeneous catalysis.
• Complex-ion formation. The same energetically accessible vacant d orbitals can accept lone pairs donated by ligands, forming dative (coordinate) bonds directly.

(Coloured compounds, the fourth property, comes from a different mechanism — d-orbital splitting — covered in Transition Elements: Colour, Stereoisomerism and Stability Constants.)

Ligands and complexes

A ligand is a species with a lone pair of electrons that forms a dative covalent bond to a central metal atom or ion. A complex is a molecule or ion formed by a central metal atom/ion surrounded by one or more ligands.

Ligands are classified by how many dative bonds each one forms to the metal:

• Monodentate (one dative bond per ligand): H₂O, NH₃, Cl⁻, CN⁻.
• Bidentate (two dative bonds per ligand): 1,2-diaminoethane (H₂NCH₂CH₂NH₂, “en”), and the ethanedioate ion (C₂O₄²⁻) — each has two separate lone-pair-bearing atoms that both bind the same metal centre.
• Polydentate (more than two dative bonds per ligand): EDTA⁴⁻, which can wrap around a metal ion using six donor atoms at once.

Coordination number is the total number of dative bonds from ligands to the central metal ion. Geometry follows from coordination number: 2 → linear; 4 → tetrahedral or square planar; 6 → octahedral.

Worked example. Cr³⁺ forms a complex with six monodentate H₂O ligands. Give the formula and charge of the complex, and its coordination number.

Six monodentate ligands means coordination number 6, so the geometry is octahedral: [Cr(H₂O)₆]³⁺ (overall charge = metal’s +3, since water is neutral).

Ligand exchange

Because the bonds between a metal ion and its ligands are dative, one ligand can be displaced by another — ligand exchange — often changing the complex’s colour and geometry. Copper(II) and cobalt(II) are the standard examples:

• [Cu(H₂O)₆]²⁺ (pale blue, octahedral) + excess NH₃(aq) → [Cu(NH₃)₄(H₂O)₂]²⁺ (deep blue, distorted octahedral) — only four of the six water ligands are replaced.
• [Cu(H₂O)₆]²⁺ (pale blue) + excess concentrated HCl → [CuCl₄]²⁻ (yellow, tetrahedral) — the change in coordination number from 6 to 4 reflects Cl⁻‘s larger size, which can’t fit six around the same copper ion.
• [Co(H₂O)₆]²⁺ (pink, octahedral) + excess NH₃(aq) → [Co(NH₃)₆]²⁺ (straw/yellow-brown, octahedral).
• [Co(H₂O)₆]²⁺ (pink) + excess concentrated HCl → [CoCl₄]²⁻ (blue, tetrahedral).

With limited (rather than excess) OH⁻ or NH₃, a precipitate forms instead of a soluble ligand-exchanged complex — e.g. Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s), a pale blue precipitate, since OH⁻ here acts as a base rather than displacing water as a ligand.

Transition-metal redox chemistry

Multiple accessible oxidation states make transition elements central to redox titrations. Feasibility of a proposed redox reaction is predicted the same way as for any other species: combine the two relevant E° values, and a positive E°_cell indicates the reaction is feasible.

MnO₄⁻/Fe²⁺ (acidified potassium manganate(VII) titration):

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

Worked example. 25.0 cm³ of Fe²⁺(aq) solution required 23.50 cm³ of 0.0200 mol dm⁻³ KMnO₄(aq) to reach the end point. Calculate the concentration of the Fe²⁺(aq) solution.

moles MnO₄⁻ = 0.02350 × 0.0200 = 4.70 × 10⁻⁴ mol

moles Fe²⁺ = 5 × moles MnO₄⁻ = 5 × 4.70 × 10⁻⁴ = 2.35 × 10⁻³ mol

[Fe²⁺] = 2.35 × 10⁻³ / 0.0250 = 0.0940 mol dm⁻³

MnO₄⁻/C₂O₄²⁻ (acidified, warm):

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

MnO₄⁻ itself is intensely purple and C₂O₄²⁻/Mn²⁺ are colourless, so this titration is famously self-indicating — the end point is the first permanent trace of pink/purple, since no more MnO₄⁻ is being consumed.

Cu²⁺/I⁻ (iodometric determination of copper):

2Cu²⁺ + 4I⁻ → 2CuI(s) + I₂

Copper(II) oxidises iodide to iodine while itself being reduced to copper(I), which immediately precipitates as white CuI. The iodine liberated is then titrated against standardised sodium thiosulfate, S₂O₃²⁻ + I₂ pathway, to determine the amount of Cu²⁺ originally present — an indirect method, since Cu²⁺ can’t be titrated directly against a reagent that gives a sharp end point on its own.

Common mistakes

Assuming every d-block element is a transition element. Scandium and zinc are the standard exceptions in the first row — check for an incomplete d sub-shell in at least one stable ion, not just d-block position.

Forgetting that ligand exchange with limited (not excess) reagent often gives a precipitate, not a soluble complex. Read the reagent volume/excess carefully in a question before deciding which outcome applies.

Losing track of geometry when the ligand changes size. Six H₂O or NH₃ ligands fit around most first-row transition ions (octahedral), but the larger Cl⁻ ligand typically only fits four (tetrahedral) — this is why [CuCl₄]²⁻ and [CoCl₄]²⁻ have a different coordination number from their aqua complexes.

Forgetting to balance the electrons transferred before combining half-equations, especially for the less familiar Cu²⁺/I⁻ system, where Cu is reduced by only one electron per ion (Cu²⁺ → Cu⁺) while I⁻ is oxidised two electrons per I₂ formed — the ×2 and ×1 multipliers needed are easy to get backwards.

Quick revision checklist

• Transition element: d-block, forms a stable ion with an incomplete d sub-shell (excludes Sc, Zn in the first row)
• Variable oxidation states, catalysis and complex formation all trace back to the closeness of 3d/4s energy and accessible vacant d orbitals
• Ligand: lone pair, dative bond to metal; mono-/bi-/polydentate by number of bonds formed
• Coordination number 4 → tetrahedral or square planar; 6 → octahedral
• Ligand exchange: excess NH₃/HCl gives a new complex; limited OH⁻/NH₃ often gives a precipitate instead
• MnO₄⁻/Fe²⁺ 1:5; MnO₄⁻/C₂O₄²⁻ 2:5 (self-indicating); Cu²⁺/I⁻ 2:4, indirect iodometric determination

Related resources

• Redox Processes: Oxidation Numbers and Electron Transfer — half-equation balancing this topic applies to transition-metal systems
• Electrochemistry: Electrolysis and Standard Electrode Potentials — E° values used to predict redox feasibility here
• Transition Elements: Colour, Stereoisomerism and Stability Constants — continues this topic’s remaining subtopics
• Cambridge AS & A Level Chemistry hub

Written against Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. Always check the current syllabus for your examination year.