Intro to Transition Metals & Complex Formation Revision Notes | A Level AQA Chemistry

1. Why Transition Metals Are Different

Scan across Period 4 of the periodic table and something unusual happens between calcium and gallium: nine elements — titanium through to copper — that behave in ways the s-block metals do not. They form coloured solutions, they catalyse reactions, and they bond to surrounding molecules in ways that produce striking, stable structures called complexes.

The reason for all of this sits in the 3d sub-level. For the s-block metals, the outer electrons are in s orbitals and that is the end of the story. For the transition metals, the 3d sub-level is being filled across the period — and it is that incomplete d sub-level that drives every characteristic property you need to know.

2. The AQA Definition: What Actually Makes a Transition Metal

Here is where many students go wrong at A Level. The AQA definition of a transition metal is not 'an element in the d-block'. The correct definition is: a transition metal is an element that forms at least one stable ion with an incomplete d sub-level.

Scandium and zinc sit at either end of the first-row d-block, but neither qualifies as a transition metal under this definition. Scandium forms only $\text{Sc}^{3+}$, which has the electron configuration [Ar] — a completely empty 3d sub-level. Zinc forms only $\text{Zn}^{2+}$, which has the configuration [Ar] $3d^{10}$ — a completely full 3d sub-level. Neither ion has an incomplete d sub-level, so neither element is a transition metal by the AQA definition.

Every element from titanium (Ti) to copper (Cu) forms at least one ion with a partially filled 3d sub-level, which is why this range is the focus of the AQA specification. Copper is a particularly interesting case — its ground-state electron configuration is [Ar] $3d^{10}$ $4s^1$ rather than the expected [Ar] $3d^9$ $4s^2$, because a full 3d sub-level is especially stable. However, copper forms $\text{Cu}^{2+}$ with configuration [Ar] $3d^9$, which is incomplete, so copper qualifies.

3. Characteristic Properties of Transition Metals

The incomplete d sub-level is not a minor detail — it is the structural reason behind four characteristic properties that AQA expects you to know and explain. These properties set transition metals apart from the s-block metals such as sodium and calcium.

• Variable oxidation states — the 3d and 4s electrons are close in energy, so different numbers of electrons can be removed to give stable ions
• Formation of coloured ions — d-d electron transitions absorb visible light when ligands split the d sub-level into two energy levels
• Catalytic activity — both as heterogeneous catalysts (e.g. iron in the Haber process) and as homogeneous catalysts (e.g. $\text{Fe}^{2+}$/$\text{Fe}^{3+}$ in the iodide–persulfate reaction)
• Complex formation — the ability to accept lone pairs from ligands into empty d orbitals, forming co-ordinate bonds

Variable oxidation states deserve particular attention. For iron, common oxidation states are +2 and +3; for manganese, the range extends from +2 all the way to +7 (as in $\text{MnO}_4^-$). This variability arises because the 3d and 4s electrons have similar energies, making it energetically feasible to remove varying numbers of electrons without the ion becoming unstable.

4. Ligands and Co-ordinate Bond Formation

Now that you have the electronic picture of transition metals, you can understand why they form complexes. A transition metal ion has available empty or partially filled d orbitals that can accept lone pairs from ligands. Certain molecules and ions have lone pairs of electrons. When a lone pair is donated into an available orbital on the metal, a co-ordinate bond forms — and that is the foundation of complex chemistry.

The key word in the definition is donating. The ligand supplies both electrons in the bond; the metal supplies the empty orbital to receive them. This is a co-ordinate bond (also called a dative covalent bond), and it is chemically identical to any other covalent bond once formed — the distinction is only in how it was made.

Ligands are classified by how many co-ordinate bonds they can form with a single metal centre. A monodentate ligand donates one lone pair and forms one co-ordinate bond — water and ammonia are the most common examples. A bidentate ligand has two donor atoms and forms two co-ordinate bonds simultaneously; ethane-1,2-diamine ($\text{H}_2\text{N}\text{CH}_2\text{CH}_2\text{NH}_2$, often abbreviated as en) and the ethanedioate ion ($\text{C}_2\text{O}_4^{2-}$) are bidentate. A polydentate ligand such as EDTA$^{4-}$ can form six co-ordinate bonds with a single metal ion.

5. Complexes and Co-ordination Number

Once ligands bond to a metal ion, the resulting species is called a complex. The AQA definition is precise: a complex is a central metal atom or ion surrounded by ligands. The overall species is often written in square brackets, for example $[\text{Cu}(\text{H}_2\text{O})_6]^{2+}$ or $[\text{CoCl}_4]^{2-}$.

The most common co-ordination numbers at A Level are 6 and 4. A co-ordination number of 6 gives an octahedral geometry, as seen in $[\text{Fe}(\text{H}_2\text{O})_6]^{3+}$. A co-ordination number of 4 gives either a tetrahedral geometry, as in $[\text{CoCl}_4]^{2-}$, or a square planar geometry, as in $[\text{Pt}(\text{NH}_3)_2\text{Cl}_2]$ (cisplatin). Square planar geometry is particularly associated with $\text{Pt}^{2+}$ and $\text{Ni}^{2+}$ complexes.

Notice that larger ligands tend to produce lower co-ordination numbers. Chloride ions ($\text{Cl}^-$) are significantly larger than water molecules, so fewer of them fit around the metal centre — this is why $[\text{CuCl}_4]^{2-}$ has a co-ordination number of 4 while $[\text{Cu}(\text{H}_2\text{O})_6]^{2+}$ has a co-ordination number of 6.

Worked example: Determining the Co-ordination Number of a Complex

State the co-ordination number of the complex $[\text{Co}(\text{en})_3]^{3+}$, where en = ethane-1,2-diamine.

1. Ligand: ethane-1,2-diamine (en). Each en molecule has two nitrogen donor atoms, so it is bidentate — it forms 2 co-ordinate bonds per ligand. — Denticity tells you how many bonds each ligand contributes. This is the step students most often skip.
2. The formula shows 3 en ligands: $[\text{Co}(\text{en})_3]^{3+}$ — The subscript 3 outside the ligand symbol tells you there are 3 en molecules surrounding the cobalt ion.
3. Co-ordination number = number of ligands × bonds per ligand = $3 \times 2 = 6$ — Co-ordination number counts bonds, not ligands. Three bidentate ligands each form 2 bonds, giving 6 co-ordinate bonds in total.
4. Co-ordination number 6 → octahedral geometry — A co-ordination number of 6 always gives an octahedral complex at A Level.

Answer: Co-ordination number = 6; geometry = octahedral

6. Naming Complexes: The Systematic Approach

AQA expects you to name common complexes and interpret names to deduce formulae. The naming rules follow a consistent pattern, and once you know the pattern, any complex becomes readable.

• Name ligands first, in alphabetical order, with a prefix indicating how many are present (di-, tri-, tetra-, penta-, hexa-). Use bis-, tris- for ligands whose names already contain a multiplying prefix.
• Name the metal second, followed by its oxidation state in Roman numerals in parentheses.
• If the overall complex ion is negative, add the suffix -ate to the metal name (e.g. cobalt becomes cobaltate, iron becomes ferrate, copper becomes cuprate).
• Common ligand names: water → aqua; ammonia → ammine; chloride → chlorido; cyanide → cyanido; hydroxide → hydroxido.

Worked example: Naming the Complex $[\text{Fe}(\text{H}_2\text{O})_6]^{3+}$

Give the systematic name of $[\text{Fe}(\text{H}_2\text{O})_6]^{3+}$.

1. Ligand: $\text{H}_2\text{O}$ → aqua. There are 6 water molecules → hexaaqua. — Always name ligands first. The prefix hexa- indicates 6 of the same ligand.
2. Metal: iron (Fe). Overall charge = +3. Each $\text{H}_2\text{O}$ is neutral (charge = 0). So: Fe oxidation state + 6(0) = +3 → Fe is +3. — Use the overall charge of the complex and the known charges of the ligands to find the metal oxidation state.
3. Overall charge = +3 (cationic), so no -ate suffix is needed. — The -ate suffix is only added when the overall complex ion carries a negative charge.
4. hexaaquairon(III) — Ligand name(s) + metal name + oxidation state in Roman numerals. No spaces between ligand prefix and metal name.

Answer: hexaaquairon(III)

7. Exam Technique and Common Mistakes

Transition metal questions on AQA papers tend to test precision of language as much as factual recall. The marking schemes are unforgiving about vague phrasing, so the following guidance is worth reading carefully before your exam.

The core thread running through this entire topic is the incomplete d sub-level. It explains why Ti–Cu are transition metals, why they show variable oxidation states, and — most directly — why they can accept lone pairs from ligands to form stable complexes. Keep that electronic picture in mind and the definitions, naming rules, and co-ordination number calculations all follow logically from it.