Transition Metals: General Properties

A transition metal is a d-block element that forms at least one ion with an incomplete (partially filled) d-subshell. This single-sentence definition does two strict pieces of work simultaneously. First, it sits the transition metals firmly inside the d-block — the ten-column central section of the periodic table spanning groups 3 to 12 — but second, and more important, it carves out two d-block elements that fail the test. Scandium (Sc) forms only Sc³⁺, which has the configuration [Ar] with an empty 3d-subshell (3d⁰); zinc (Zn) forms only Zn²⁺, with a completely full 3d¹⁰. Neither ion has a partially filled d-subshell. They are d-block, but they are not transition metals in the strict AQA sense. This distinction matters because the characteristic chemistry of the transition metals — variable oxidation states, complex-ion formation with ligands, coloured compounds and solutions, paramagnetism, and prolific catalytic activity (both homogeneous and heterogeneous) — flows directly from having unpaired d-electrons available to participate in bonding and absorption of visible light. This lesson establishes the definition rigorously, audits the electron configurations of the first-row transition series (with the Cr and Cu anomalies), explains why 4s is lost before 3d on ionisation, surveys the five characteristic properties, and previews the deeper treatments in lessons 5–9.

Spec mapping (AQA 7405): This lesson anchors §3.2.5 (transition metals — general properties of transition elements). Complex-ion formation is developed in lesson 5; ligand substitution reactions in lesson 6; the origin of colour and d–d transitions in lesson 7; aqueous-ion chemistry (acidity, hydrolysis, hydroxide precipitation) in lesson 8; homogeneous and heterogeneous catalysis in lesson 9. The electron-configuration material here also reinforces §3.1.1.3 (electron configurations of atoms and ions, including the anomalies of Cr and Cu). Refer to the official AQA specification document for the exact wording of each section.

Assessment objectives: AO1 marks reward the strict definition of a transition metal, recall of the characteristic properties (variable oxidation states; complex-ion formation; coloured compounds; catalysis; paramagnetism), and recall of the common oxidation states of Sc–Zn. AO2 marks reward writing the full electron configurations of Sc–Zn atoms (including the Cr and Cu anomalies) and the corresponding M^n+ ions (4s lost first), and predicting which ions a given element can sensibly form. AO3 marks reward rationalising why Sc and Zn fail the transition-metal test, why Cr and Cu adopt their anomalous ground-state configurations (half-full and full d-subshell stability via exchange energy), and why the characteristic properties cluster together (they all derive from the partially filled d-subshell).

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Definition — The Strict AQA Wording

Key Definition: A transition metal is a d-block element that forms at least one ion with an incomplete d-subshell.

Read this slowly. Three clauses, each load-bearing:

1. "d-block element" — the element's last-filling electrons go into a 3d (or 4d, 5d) subshell. For Period 4 this means Sc–Zn (groups 3–12 inclusive).
2. "at least one ion" — any stable ion of the element qualifies; the element does not have to form an incomplete-d ion in every compound, only in some. Cu satisfies this via Cu²⁺ (3d⁹) even though Cu⁺ (3d¹⁰) does not.
3. "incomplete d-subshell" — strictly between 1 and 9 d-electrons inclusive. Neither 3d⁰ nor 3d¹⁰ counts.

What this excludes

• Scandium (Sc): atomic configuration [Ar] 3d¹ 4s². The only common ion is Sc³⁺, configuration [Ar] (3d⁰). The d-subshell is empty, not partially filled. Sc is therefore not a transition metal in the AQA sense, even though it sits in the d-block and is sometimes loosely included in older textbooks.
• Zinc (Zn): atomic configuration [Ar] 3d¹⁰ 4s². The only common ion is Zn²⁺, configuration [Ar] 3d¹⁰. The d-subshell is completely full. Zn is therefore not a transition metal either.

These exclusions are not pedantry. They predict observable chemistry:

• Sc³⁺ and Zn²⁺ solutions are colourless — no d–d transitions are possible (no electrons to promote, in Sc³⁺'s case; no empty d-orbital to promote into, in Zn²⁺'s case). Compare with Cu²⁺ (blue), Fe³⁺ (yellow–brown), [Co(H₂O)₆]²⁺ (pink).
• Sc and Zn show only one stable oxidation state in compounds (Sc +3; Zn +2). They do not exhibit the variable-oxidation-state hallmark of the true transition metals.
• Sc and Zn compounds are generally not good catalysts; their salts do not display the rich redox chemistry of, say, Fe²⁺/Fe³⁺ or Mn²⁺/MnO₄⁻.

Exam Tip: "d-block element" and "transition metal" are not synonyms. Sc and Zn are d-block but not transition. A question that asks you to "explain why scandium is not a transition metal" wants the strict definition applied to Sc³⁺'s 3d⁰ configuration — not a vague comment about reactivity.

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Electron Configurations of Sc–Zn

The atoms

The 4s subshell fills before 3d for neutral atoms of Sc through Zn, because in the unionised atom the 4s orbital is slightly lower in energy than the 3d orbitals. The ten Period 4 d-block atoms are:

Z	Element	Symbol	Atomic configuration
21	Scandium	Sc	[Ar] 3d¹ 4s²
22	Titanium	Ti	[Ar] 3d² 4s²
23	Vanadium	V	[Ar] 3d³ 4s²
24	Chromium	Cr	[Ar] 3d⁵ 4s¹     (anomaly)
25	Manganese	Mn	[Ar] 3d⁵ 4s²
26	Iron	Fe	[Ar] 3d⁶ 4s²
27	Cobalt	Co	[Ar] 3d⁷ 4s²
28	Nickel	Ni	[Ar] 3d⁸ 4s²
29	Copper	Cu	[Ar] 3d¹⁰ 4s¹    (anomaly)
30	Zinc	Zn	[Ar] 3d¹⁰ 4s²

Eight of these ten configurations follow the predictable pattern "fill 4s, then add electrons to 3d one at a time, following Hund's rule". Two do not.

The Cr and Cu anomalies — exchange energy

Chromium could in principle be [Ar] 3d⁴ 4s² by simple aufbau. In fact it is [Ar] 3d⁵ 4s¹ — one 4s electron has been "promoted" into the 3d set so that all five 3d orbitals are singly occupied with parallel spins. The driver is exchange energy: when electrons in different orbitals of the same subshell have parallel spins, a quantum-mechanical stabilisation arises from the indistinguishability of identical-spin electrons. Each pair of parallel-spin electrons in the same subshell contributes one "exchange term" K to the total energy. For a half-full d⁵ high-spin configuration there are C(5,2) = 10 parallel pairs and thus 10K of exchange stabilisation — a maximum for any single subshell. The energy cost of promoting one electron 4s→3d is more than repaid by the larger exchange-stabilisation gain.

Copper could in principle be [Ar] 3d⁹ 4s². In fact it is [Ar] 3d¹⁰ 4s¹ — again one 4s electron has moved into 3d, this time completing the d-subshell. The driver is similar in spirit: a fully filled 3d¹⁰ subshell is especially stable (fully occupied with paired spins; closed-subshell symmetry), and the relatively small 4s→3d energy cost is repaid by the additional pairing and shielding gains.

Common Misconception: The "extra stability of half-filled and full d-subshells" is often stated as a slogan without explanation. The real driver is exchange energy (parallel-spin stabilisation) for Cr, and a combination of exchange and closed-subshell stability for Cu. At A* you should be able to name the mechanism, not just the result.

The ions — 4s lost before 3d

Once you start removing electrons to form a cation, the 4s electrons are lost first, then the 3d. The textbook way of justifying this is that, once the atom is partly ionised, the 4s orbital lies above the 3d set in energy — the 3d orbitals are more strongly attracted to the nucleus once the 4s shielding electrons have been removed. (At full A* level the picture is more subtle: the rule "4s fills first but ionises first" reflects the relative energies of the orbitals in the neutral vs ionised environment, and is best treated as an empirical rule rather than a rigid theorem. The Going Further section says more.)

Apply the rule mechanically: write the atom's configuration, strip the 4s electrons, then any 3d electrons needed to reach the required charge.

Ion	Starting atom	Procedure	Final configuration	d-count
Sc³⁺	[Ar] 3d¹ 4s²	lose 2 × 4s, then 1 × 3d	[Ar]	0
Ti²⁺	[Ar] 3d² 4s²	lose 2 × 4s	[Ar] 3d²	2
Ti⁴⁺	[Ar] 3d² 4s²	lose 2 × 4s + 2 × 3d	[Ar]	0
V³⁺	[Ar] 3d³ 4s²	lose 2 × 4s + 1 × 3d	[Ar] 3d²	2
Cr³⁺	[Ar] 3d⁵ 4s¹	lose 1 × 4s + 2 × 3d	[Ar] 3d³	3
Mn²⁺	[Ar] 3d⁵ 4s²	lose 2 × 4s	[Ar] 3d⁵	5
Fe²⁺	[Ar] 3d⁶ 4s²	lose 2 × 4s	[Ar] 3d⁶	6
Fe³⁺	[Ar] 3d⁶ 4s²	lose 2 × 4s + 1 × 3d	[Ar] 3d⁵	5
Co²⁺	[Ar] 3d⁷ 4s²	lose 2 × 4s	[Ar] 3d⁷	7
Ni²⁺	[Ar] 3d⁸ 4s²	lose 2 × 4s	[Ar] 3d⁸	8
Cu⁺	[Ar] 3d¹⁰ 4s¹	lose 1 × 4s	[Ar] 3d¹⁰	10
Cu²⁺	[Ar] 3d¹⁰ 4s¹	lose 1 × 4s + 1 × 3d	[Ar] 3d⁹	9
Zn²⁺	[Ar] 3d¹⁰ 4s²	lose 2 × 4s	[Ar] 3d¹⁰	10

Two observations from this table you should be able to make under exam pressure:

• Fe³⁺ (3d⁵) is unusually stable. The half-full d-subshell brings the same maximal exchange stabilisation that made Cr atomic [Ar] 3d⁵ 4s¹ favourable. This is part of why Fe²⁺ is readily oxidised to Fe³⁺ in aerated solutions.
• Sc³⁺ is 3d⁰ and Zn²⁺ is 3d¹⁰. Both ions formally complete the [Ar] or [Ar]3d¹⁰ closed-shell structure — hence the very high stability of these single oxidation states and the absence of variable-ON chemistry.

Worked example — MnO₄⁻. What is the oxidation number of manganese in MnO₄⁻, and what is the d-electron count? Oxygen is −2; four oxygens contribute −8; the overall charge is −1; therefore Mn must be +7. Atomic Mn is [Ar] 3d⁵ 4s². Strip 2 × 4s and 5 × 3d → [Ar], i.e. Mn(VII) is d⁰. This is why MnO₄⁻'s intense purple colour is not a d–d transition but a charge-transfer transition (LMCT, ligand-to-metal charge transfer — see lesson 7).

Common Misconception: Students frequently write Fe²⁺ as [Ar] 3d⁴ 4s² — removing electrons from 3d first. Wrong. The 4s electrons go first. The correct configuration is [Ar] 3d⁶. This trips students up because it contradicts the "4s fills first" aufbau rule they learned for atoms, but the orbital ordering inverts once ionisation begins.

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The Five Characteristic Properties

The transition metals share five clusters of properties that, taken together, set them apart from the s-block and p-block. Each is rooted in the partially filled d-subshell.

1. Variable oxidation states

Almost every transition metal exhibits multiple stable oxidation states in its compounds. The reason is that 3d and 4s lie close enough in energy that successive ionisations remove electrons one (or a few) at a time without prohibitive energy cost — unlike, say, sodium, where removing the second electron requires breaking into the inert [Ne] core. Common AQA-table oxidation states for Period 4:

Element	Common oxidation states	Most stable	Notable feature
Ti	+2, +3, +4	+4	d⁰ in TiO₂
V	+2, +3, +4, +5	+5	All four observable in aqueous solution (lesson 8)
Cr	+2, +3, +6	+3	Cr(VI) = strong oxidant (CrO₄²⁻, Cr₂O₇²⁻)
Mn	+2, +3, +4, +6, +7	+2	Widest spread; +7 in MnO₄⁻
Fe	+2, +3	+3	Half-full d⁵ stability favours Fe(III)
Co	+2, +3	+2	Co(III) only with strong ligands
Ni	+2	+2	Effectively single ON in simple chemistry
Cu	+1, +2	+2	Cu(I) disproportionates in water

Manganese, with eight possible oxidation states between +2 and +7 (and rare +1 in some complexes), is the species with the widest range across the row — a direct consequence of having a half-filled d⁵ atomic configuration giving it the most electrons to "play with" while still preserving stability.

2. Complex-ion formation

Transition-metal ions are powerful Lewis acids: small radius, high charge density, and (critically) empty 3d, 4s, and 4p orbitals available to accept lone pairs from electron-donor species called ligands. The resulting M–L bonds are dative covalent (coordinate) bonds — both shared electrons come from the ligand. Common ligand types include H₂O, NH₃, Cl⁻, OH⁻, CN⁻; common geometries are octahedral (6 ligands, e.g. [Cu(H₂O)₆]²⁺, [Co(H₂O)₆]²⁺, [Fe(H₂O)₆]³⁺), tetrahedral (4 ligands, e.g. [CuCl₄]²⁻), and square planar (4 ligands in a special geometry, e.g. cisplatin [PtCl₂(NH₃)₂]). Lesson 5 develops complex-ion structure and naming in full.

3. Coloured compounds and solutions

Aqueous solutions of transition-metal ions are typically vivid: Cu²⁺ blue, Fe³⁺ yellow–brown, Co²⁺ pink, Ni²⁺ green, Cr³⁺ green, MnO₄⁻ intense purple. The mechanism is d–d electronic transitions: in the presence of ligands, the five 3d orbitals split into two sets at slightly different energies (octahedral splitting: lower-energy t₂_g triplet and higher-energy e_g doublet; see lesson 7 for the full crystal-field picture). Visible-wavelength photons matching the splitting energy ΔE are absorbed, and the colour we perceive is the complementary colour of the absorbed band. Sc³⁺ (d⁰) and Zn²⁺ (d¹⁰) are colourless because no d–d transition is possible — there are no electrons to promote (Sc³⁺) or no empty d-orbital to promote into (Zn²⁺). This is direct experimental evidence that they fail the transition-metal test.

Practical-skills box — watching oxidation-state changes by colour. Many redox titrations in transition-metal chemistry are self-indicating: the reagent or product is intensely coloured, and the endpoint is the persistence of one drop of colour past the stoichiometric point. The classical demonstration is the iron(II) titration with potassium manganate(VII): Fe²⁺ (pale green) is oxidised to Fe³⁺ (yellow–brown) by MnO₄⁻ (purple), which is reduced to Mn²⁺ (almost colourless). Until all the Fe²⁺ is oxidised, every drop of purple MnO₄⁻ is decolourised on entering the flask. The first persistent pink colour signals the endpoint. The same colour-change principle is exploited in the vanadium oxidation-state series (V²⁺ violet → V³⁺ green → VO²⁺ blue → VO₂⁺ yellow), used in Required Practical 11 to qualitatively track redox titrations. Always record the exact colour transition you observe, including the intermediate, not just initial and final.

4. Catalysis

Transition metals — elemental, as salts, or in complexes — are spectacularly effective catalysts. The two underlying capabilities are (a) variable oxidation states (the metal can cycle between two oxidation states during the catalytic cycle, accepting electrons from one reactant and donating them to another, lowering activation energy), and (b) the ability to form complexes with the reactants (binding reactant molecules in close proximity and correct orientation). Industrial examples to know:

Catalyst	Process	Reaction	Type
Fe	Haber process	N₂ + 3H₂ ⇌ 2NH₃	Heterogeneous
V₂O₅	Contact process	2SO₂ + O₂ ⇌ 2SO₃	Heterogeneous
Ni	Hydrogenation of alkenes	C=C + H₂ → C–C	Heterogeneous
MnO₂	Decomposition of H₂O₂	2H₂O₂ → 2H₂O + O₂	Heterogeneous
Pt / Pd / Rh	Catalytic converter	2CO + 2NO → 2CO₂ + N₂	Heterogeneous
Fe²⁺ / Fe³⁺	S₂O₈²⁻ + I⁻ reaction	(lesson 9)	Homogeneous

Catalysis is developed properly in lesson 9.

5. Magnetic properties (paramagnetism)

Any species with unpaired electrons is paramagnetic: it is weakly attracted into a magnetic field. Transition-metal ions with partially filled d-subshells almost always have unpaired d-electrons (Hund's rule), so they are typically paramagnetic. The number of unpaired electrons is predictable from the d-count and the type of ligand field (see lesson 7 for high-spin vs low-spin). A simple correlation: paramagnetic moments increase with the number of unpaired electrons. Sc³⁺ (d⁰) and Zn²⁺ (d¹⁰) are diamagnetic — zero unpaired electrons — again consistent with them not being transition metals in the strict sense.

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Worked Examples — Writing Configurations

Example 1 — Fe²⁺.

Atomic Fe is [Ar] 3d⁶ 4s². Remove two electrons; 4s goes first; both 4s electrons removed. Fe²⁺ = [Ar] 3d⁶. Six d-electrons; four unpaired by Hund's rule in the high-spin case (typical for weak-field aqua complexes).

Example 2 — Fe³⁺.

Atomic Fe is [Ar] 3d⁶ 4s². Remove three electrons; 4s first (−2), then 3d (−1). Fe³⁺ = [Ar] 3d⁵. Half-full — maximal exchange stabilisation — hence Fe³⁺ is thermodynamically favoured over Fe²⁺ in aerated aqueous solution.

Example 3 — Cu²⁺.

Atomic Cu is [Ar] 3d¹⁰ 4s¹ (anomaly). Remove two electrons; 4s first (−1), then 3d (−1). Cu²⁺ = [Ar] 3d⁹. Nine d-electrons; one unpaired; the d¹ hole drives the intense blue colour of [Cu(H₂O)₆]²⁺.

Example 4 — MnO₄⁻ (manganate(VII)).